Solutions and Solubility

Solutions and Solubility
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
1. Temperature
For solid solutes: as temperature
increases, solubility increases.
For gas solutes: as temperature
increases, solubility decreases.
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.”
Solute/Solvents cont.
Solubility Summary
Solute Type
Nonpolar (Fat
Grease)
Polar
Ionic (salt)
High solubility-soluble
Nonpolar
solvent
Soluble (soap)
Insoluble
insoluble
Polar Solvent
Insoluble
(water)
Soluble (water)
Soluble (water)
Low solubility-insoluble
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
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.
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
Examples
•
•
•
•
•
•
AgNO3 = water soluble
AgCl = insoluble
Na2S = soluble
NaCl = soluble
CaCO3 = insoluble
AlPO4 = insoluble
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.
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.
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.
Molarity
• Molarity (M)- number of moles of solute in 1L of
solution. Table T
Molarity= moles of solute
liters of solution
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
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
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
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
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)
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.
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