1

Ions in Aqueous Solutions and Colligative Properties

Chapter 14

2

Dissociation

• The separation of ions that occurs when an ionic compound dissolves.

3

Dissociation examples

ClNaNaCl

ClCaCaCl 222

4

You try

• Write the equation for the dissolution of NH4NO3 in water. If 1 mol of ammonium nitrate is dissolved, how many moles of each type of ion are produced?

• 1 mol of each type of ion

3434 NONHNONH

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Precipitation Reactions

• When two solutions are mixed, a double replacement reaction may occur.

• If one of the products is insoluble, it will form a precipitate.

• See table 14-1 on page 427

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Example

• Solutions of (NH4)2S and Cd(NO3)2 are mixed. Will a precipitate form?

2424 2)( SNHSNH

32

23 2)( NOCdNOCd

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Example continued

CdSNONHNOCdSNH 342324 2)(

)()(2

)()()(

34

2324

sCdSaqNONH

aqNOCdaqSNH

The cadmium sulfide is the precipitate.

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Net Ionic Equations

• Includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution.

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Spectator Ions

• Ions that do not take part in a chemical reaction and are found in solution both before and after the reaction.

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Example

)(2)(2)(

)()(2)(2)(

43

243

2

aqNHaqNOsCdS

aqSaqNHaqNOaqCd

)()()( 22 sCdSaqSaqCd

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You try

• A solution of sodium sulfide is combined with a solution of iron(II) nitrate. Does a precipitate form?

• Iron(II) sulfide is the precipitate.

)()(2

)()()(

3

232

sFeSaqNaNO

aqNOFeaqSNa

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You try continued

• Write the net ionic equation for the previous reaction.

)()(2)(2

)(2)()()(2

3

322

sFeSaqNOaqNa

aqNOaqFeaqSaqNa

)()()( 22 sFeSaqFeaqS

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Ionization

• The process that forms ions from solute molecules by the action of the solvent.

• The attraction between the solvent and the solute is strong enough to break the covalent bonds.

)()( aqClaqHHCl

14

Hydronium

• H3O+

• Formed when an H+ ion is combined with a water molecule (hydrated).

• Happens instantly when H+ ions are in water.

• Highly exothermic• Formed by many molecular compounds that

ionize

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A more accurate picture

)()( aqClaqHHCl

)()()()( 32 aqClaqOHgHCllOH

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Strong electrolytes

• Any compound whose dilute aqueous solutions conduct electricity well.

• All or almost all dissolved compound is in the form of ions– Not all compound has to dissolve, but the part

that does must be ions

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Weak electrolytes

• Any compound whose dilute aqueous solutions conduct electricity poorly.

• A small amount of the dissolved compound is in the form of ions.

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Be careful!

• Strong electrolytes have a high degree of ionization or dissociation, regardless of their concentration.

• Weak electrolytes have a low degree of ionization or dissociation, regardless of their concentration.

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Colligative Properties

• Properties of solutions that depend on the concentration of solute particles, but not the identity of solute particles.

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Nonvolatile substance

• Has little tendency to become a gas under existing conditions.

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Vapor-pressure lowering

• The vapor pressure of a solvent containing a nonvolatile solute is lower than the vapor pressure of the pure solvent at the same temperature.

• The solute lowers the concentration of solvent molecules at the surface.

• Fewer molecules enter the vapor phase.

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Effects

• See figure 14-6 on page 436

• The solution remains liquid over a wider temperature range.

• The freezing point is lowered and the boiling point is raised.

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Molal freezing-point constant

• Kf

• The freezing point depression of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute.

• = -1.86 °C/m for water

• 2 molal decreases 3.72 °C

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Freezing-point depression

• The difference between the freezing points of the pure solvent and a solution of a nonelectrolyte in that solvent.

• It is directly proportional to the molal concentration of the solution.

mKt ff

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Example

• Determine the freezing point of a water solution of fructose, C6H12O6 made by dissolving 58.0 g of fructose in 185 g of water.

• -3.24 °C

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You try

• Determine the molal concentration of a solution of ethylene glycol, HOCH2CH2OH, if the solution’s freezing point is -6.40 °C.

• 3.44 m

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Molal boiling-point constant

• The boiling-point elevation of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute.

• Kb = 0.51 °C/m for water

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Boiling-point elevation

• The difference between the boiling points of the pure solvent and a nonelectrolyte solution of that solvent.

• Directly proportional to the molal concentration of the solution

mKt bb

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Example

• What is the boiling point of a solution of 25.0 g of 2-butoxyethanol, HOCH2CH2OC4H9, in 68.7 g of ether?

• 40.8 °C

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You try

• What mass of glycerol, CH2OHCHOHCH2OH, must be dissolved in 1.00 kg of water in order to have a boiling point of 104.5 °C?

• 810 g

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Semipermeable membrane

• Allows the movement of some particles while blocking the movement of others.

• Example: allows water molecules through, but not sucrose molecules

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Osmosis

• The movement of solvent through a semipermeable membrane from the side of lower solute concentration to the side of higher solute concentration.

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Osmotic pressure

• The external pressure that must be applied to stop osmosis.

• The greater the concentration of a solution, the greater the osmotic pressure.

35

Electrolytes

• 1 mole of an electrolyte produces more than one mole of particles in solution.– The ions separate

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Example

• A water solution contains 42.9 g of calcium nitrate dissolved in 500. g of water. Calculate the freezing point of the solution.

• -2.92 °C

37

You try

• What is the expected boiling point of a 1.70 m solution of sodium sulfate in water?

• 102.6 °C

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Actual values

• Our expected values are not always what is observed.

• See table 14-3 on page 445

• Differences are caused by attractive forces between ions in solution.