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Solutions and Solubility

This article discusses the properties of solutions and solubility, including temperature and pressure effects, solubility curves, and solubility rules. It also covers precipitation reactions, concentrations of solutions, parts per million, and colligative properties such as freezing and boiling point changes.

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Solutions and Solubility

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  1. Solutions and Solubility

  2. Solutions • Solution – a homogenous mixture of a solute dissolved in a solvent. The solubility (ability to dissolve) of a solute in a solvent is dependent on the • Temperature For solid solutes: as temperature increases, solubility increases. For gas solutes: as temperature increases, solubility decreases.

  3. Solutions con 2. Pressure For solid solutes: as pressure increases, solubility remains the same. For gas solutes: as pressure increases, solubility increases 3. Nature of Solute/Solvent “Like dissolves in like.”

  4. Solute/Solvents cont. Solubility Summary High solubility-soluble Low solubility-insoluble

  5. Solubility Curves • Shows the number of grams of solute that can be dissolved in 100.g of water at temperatures between 0 degrees C and 100 degrees C. • Each line represents the maximum amount of that substance that can be dissolved at a given temperature. • Lines that show an increase in solubility as temperatures increase represent solids being dissolved in water. • Lines that show a decrease in solubility as temperatures increase represent gases being dissolved in water. These are NH3, SO2, and HCl

  6. There are three types of solutions • 1. An unsaturated solution is a solution in which more solute can be dissolved at a given temperature. • 2. a saturated solution is a solution containing the maximum amount of solute that will dissolve at a given temperature. • 3. a supersaturated solution is a solution that contains more solute than would dissolve in a saturated solution at a given temperature.

  7. Solubility Rules • Not all ionic compounds are water soluble • There are some general rules for compounds that are water soluble: • Group 1 ionic cmpds and ammonium (NH4+) are always water soluble • Group 17 ionic cmpds are water soluble except when paired with Ag, Pb, and Hg ions • See Table F for full rules and exceptions

  8. Examples • AgNO3 = water soluble • AgCl = insoluble • Na2S = soluble • NaCl = soluble • CaCO3 = insoluble • AlPO4 = insoluble

  9. Precipitation reactions • Recall that double-replacement reactions have the general formula: AB(aq) + CD(aq) AD(aq) +CB(s) A precipitation reaction will take place if one or both of the products is listed as an insoluble solid.

  10. 2KI(aq) + Pb(NO3)2(aq) -> 2KNO3(aq) + PbI(s) • KI = soluble • Pb(NO3)2 = soluble • 2KNO3(aq) = soluble • PbI(s) = insoluble solid • In a precipitation reaction, two clear aqueous solutions are combined to form a cloudy, solid precipitate that can be collected by filtration.

  11. Concentrations of Solutions • Because solutions are homogeneous mixtures, their compositions can vary. Sometimes it is adequate to refer to a solution as dilute or concentrated. These are qualitative descriptions of concentration. It is more precise to describe the concentration of solutions in quantitative measures.

  12. Molarity • Molarity (M)- number of moles of solute in 1L of solution. Table T Molarity= moles of solute liters of solution

  13. Calculating molarity Highly concentrated HCl(aq) has a molarity of 12M This means there are 12 moles of HCl dissolved in 1 Liter of water 12M = 12 moles 1 Liter

  14. Sample Problem • How many grams of NaCl must be added to 1 Liter of water to make a 3M solution? • 3M = 3 moles NaCl I Liter 1 mole NaCl = 58g x 3moles = 174g NaCl

  15. Parts Per Million • Parts per million is another way of measuring the concentration of a solution • The general formula is: Parts per million = grams of solute x 1,000,000 grams of solution

  16. Parts per million example • 5 grams of NaCl is dissolved in 2.5L of water. What is the concentration of NaCl in parts per million (ppm)? • Remember 1mL water = 1 g ppm = 5g NaCl x 1,000,000 = 2000 ppm 2500 g H2O

  17. Colligative Properties Freezing and boiling points of water change when salts (nonvolatile solutes) are added. Colligative properties depend on the number of particles in a substance 1. Freezing Point Depression: when any salt is added to water, the freezing point of the water decreases. Freezing point of pure water= 0 degrees C Freezing point of salt water (NaCl solution)= -21°C (-6°F) under controlled lab conditions In the real world, on a real sidewalk, sodium chloride can melt ice only down to about -9°C (15°F)

  18. Molecular vs. Ionic • When one mole of sucrose is dissolved in water, one mole of particle is produced in solution: C12H22O11(s) C12H22O11(aq) • When one mole of an ionic substance is dissolved in water, the results are different. The ionic substance dissociates into individual ions: NaCl(s) Na(aq) + Cl (aq) • The greater number of ions, the lower the freezing point.

  19. Boiling Point Elevation • When any salt is added to water, the boiling point of the water increases. • Boiling point of pure water= 100° C • Boiling point of salt water solution increases

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