Unit 2: Liquids & solids, solubility,
equilibrium
By: Ali Montgomery and Sam Block
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General Properties of Aqueous Solutions
• A substance such as NaCl whose aqueous solutions contain
ions is an electrolyte
• A substance such as C12H22O11 that does not form ions in
solution is a nonelectrolyte. Most molecular substances are
nonelectrolytes
• Strong electrolytes are solutes that exist in solution almost
entirely as ions
• Weak electrolytes are solutes that exist in solution mostly as
molecules with only a small fraction in the form of ions
Precipitation Reactions
• Reactions resulting in the formation of an insoluble product
are called precipitation reactions
• A precipitate is an insoluble solid formed by a reaction in
solution
• The solubility of a substance is the amount of substance that
can be dissolved in a given quantity of solvent at a specific
temperature
• Any substance with a solubility of less than 0.01 mol/L is
generally considered insoluble
http://www.iun.edu/~cpanhd/
C101webnotes/chemical
reactions/images/agcl.jpg
Solubility Rules (Appendix 2)
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http://dist113.org/dhs/Depts/Science/Hinton/appendix/a02.pdf
Rule 1 ALKALI METALS and
AMMONIUM
Most alkali metal and ammonium salts are soluble.
All common salts of hydrogen, sodium, potassium,
and ammonium are soluble.
Rule 2 ACETATES All acetates are soluble. Silver acetate is only
moderately soluble.
Rule 3 CARBONATES Carbonates are insoluble except those
covered by
Rule 1. All are soluble in the presence of acid.
Rule 4 CHLORIDES,
BROMIDES, and
IODIDES
Most metal chlorides, bromides and iodides are
soluble except those of silver and lead. Lead (II)
chloride is moderately soluble.
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Solubility Rules (continued)
Rule 5 CHROMATES Chromates are insoluble except for those covered
by Rule 1 and those of manganese (II) and iron (III).
Chromates of mercury (II), calcium, and strontium
are slightly soluble.
Rule 6 HYDROXIDES Hydroxides are insoluble except those covered by
Rule 1 and those of calcium, strontium, and barium.
Magnesium hydroxide is slightly soluble.
Rule 7 NITRATES All nitrates are soluble.
Rule 8 OXIDES Oxides are insoluble except those covered by Rule 1
and those of calcium, strontium, and barium.
Oxides that are soluble usually react with water.
Rule 9 PERMANGANATES Most permanganates are soluble.
Rule 10 PHOSPHATES Phosphates are insoluble except those covered by
Rule 1.
Rule 11 SULFATES Sulfates are soluble except those of strontium,
barium, and lead. Calcium and silver sulfates are
only moderately soluble.
Rule 12 SULFIDES Sulfides are insoluble except those covered by
Rule 1.
Ionic Equations
• Molecular Equations show the chemical formulas of reactants
and products without indicating their ionic character
• Ex. 3AgNO3(aq)
+ AlCl3(aq)  3AgCl(s) +
Al(NO3)3(aq)
• Complete ionic equations show all souble strong electrolytes
as ions
• Ex. 3Ag+
(aq) + 3NO3-(aq) + Al^3+(aq) + 3Cl-(aq)
3AgCl(s) + Al^3+(aq) + 3NO3-(aq)
• Net ionic equations do not include spectator ions, which are
quantities that can be canceled out
• Ex. 3Ag+
(aq) + 3Cl-(aq) 3AgCl(s)
Equilibrium
• Chemical Equilibrium occurs when opposing reactions
proceed at equal rates
• Although concentrations do not change at equilibrium, the
reaction still occurs
• For any equilibrium equation
aA + bB <--> dD + eE, the equilibrium can be expressed by an
equilibrium-constant expression as
• The equilibrium constant expression depends only upon the
stoichiometry of the reaction, not on its mechanism
• Solids and liquids are only 1, and do not go into the constant
because their concentrations don’t change
• For an equilibrium constant expression in terms of pressure,
Kp, use Kp=Kc(RT)^(Δn) to convert from the Kc constant
• Δn is the moles of product-moles of reactant
More with Equilibrium Constants
• If K>1, equilibrium lies to the right, and products are
favored
• If K<1, equilibrium lies to the left, and reactants are
favored
• The equilibrium-constant expression and the
equilibrium constant of the reverse of a reaction are
the reciprocals of those of the forward reaction
• When a reaction is the sum of two or more reactions,
its equilibrium constant is the product of the
equilibrium constants for the individual reactions
• Equilibrium constants vary with temperature.
I C E
nitial
hange
quilibrium
!
• If the equilibrium concentration of at least one species
involved in a chemical reaction is known, we can use
stoichiometry to calculate the other concentrations with an
ICE chart. The steps to do this are:
• 1) Put all known initial and equilibrium concentrations into an
ice chart
• 2) For species where both initial and equilibrium
concentrations are provided, calculate the change in
concentration
• 3) Use stoichiometry of the reaction (keeping in mind the
effects of coefficients on the species involved) to calculate the
change in concentration for the other species involved
• 4) Use the change in concentration to arrive at the
equilibrium concentrations.
http://images.google.com/imgres?imgurl=
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Reaction Quotients
• If the Reaction Quotient Q (a number obtained by
substituting actual product and reactant
concentrations into an equilibrium-constant
equation) equals K, the system is at equilibrium
• If Q>K, the concentration of products is too big so
the reaction shifts left to offset it by increasing the
concentration of reactants
• If Q<K, the concentration of reactants is too big, so
the reaction shifts right, which increases the amount
of products
Le Chatelier’s Principle
• Changes in temperature, pressure (or volume), or the
concentration of one substance involved in the
reaction shift the equilibrium position to offset the
changes; catalysts do not shift equilibrium
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Le Chatelier’s Principle
• Adding a substance causes the reaction to shift as more of
that excess substance is consumed. Removing a substance
causes the reaction to shift to produce more that substance
• Reducing the volume (increasing the pressure) of an
equilibrium with gaseous components causes the system to
shift to the side that reduces the number of moles; decreasing
the pressure (increasing volume) shifts the system to the side
that increases the number of moles of gas
• If the system is exothermic, heat is treated as a product.
Raising the temperature shifts the system left
• If the system is endothermic, heat is a reactant. Raising the
temperature shifts the system right
Solubility Product
• The solubility product, Ksp, equals the product
of the concentration of the ions involved in
the equilibrium, raised to the power of the
coefficient in the equation
• BaSO4 (s) <-> Ba ^2+ (aq) + SO4 ^2- (aq)
• Ksp=[Ba^2+][SO4 ^2-]
• The Ksp can be used to calculate the solubility
of a compound
Factors that Affect Solubility
• The presence of common ions reduces solubilty
• Temperature affects solubility of substances
dissolved in water
• The solubility of compounds containing basic anions
increases as the solution is made more acidic, and pH
decreases
• Salts with anions of negligible basicity, anions of
strong acids such as Cl-, are not affected by pH
changes
Precipitation
• If Q>Ksp, precipitation occurs until Q=Ksp
• If Q=Ksp, equilibrium exists in a saturated
solution
• If Q<Ksp, solid dissolves until Q=Ksp
Intermolecular Forces
• Gas molecules undergo constant, chaotic
motion
• Liquid molecules are free to move, but kept in
close proximity
• Solid molecules have strong enough attractive
forces to restrain molecular motion
IMFA’s Continued
• Dipole-dipole forces
– Dispersion force depends on polarizability, size,
and shape
• London dispersion forces
– All molecules
– Especially when molecules very close together
• Hydrogen bonding
– exist between a hydrogen atom in a polar bond
and an electronegative element, (H-O, H-F, H-N)
Flow
• Viscosity: Resistance to flow!
– Greater weight, attraction increased viscosity
• Surface tension: Energy required to increase
the surface area of a liquid by a unit amount
– Adhesive: Bind substance to surface
– Cohesive: Bind similar molecules together
• Capillary action: Rise of liquids up narrow tube
Phase Changes
In J, must convert to kJ
(vaporization)
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Phase Diagram
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Crystalline Structures
• Simple cubic: 1 atom, V=8r3, e=2r
• Body-centered: 2 atoms, V=(4r/√3)3, e=4r/√3
• Face-centered: 4 atoms, V= (32r3/√2), e=4r/√2
http://www.substech.com/
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James Bonding in Solids
• Molecular solids: atoms/molecules held together by
intermolecular forces (London dispersion, dipoledipole, hydrogen bonds), soft, relatively low boiling
points, poor thermal and electrical conduction, ex:
methane, sucrose, and dry ice
• Covalent-network solids: atoms held together by
covalent bonds, very hard, high melting points, poor
thermal and electrical conductors, ex: diamonds
• Ionic solids: ions held together by ionic bonds, hard
and brittle, high melting points, poor thermal and
electrical conduction, ex: salts
• Metallic solids: metal atoms held together by metallic
bonds, vary in strength of bonding, wide range of
physical properties (hardness, melting points),
malleable and ductile, excellent thermal and electrical
conductors, ex: copper, iron, aluminum
Saturated Solutions and Solubility
• Saturated: Solution that is in equilibrium with
undissolved solute
• Solubilty: Amount of solute needed to form a
saturated solution in a given amount of solvent
• Unsaturated: Less solute than is needed to form
a saturated solution
• Superaturated: Not unsaturated. The opposite, in
fact
Henry’s Law (It’s French)
• Sg=kPg
• Sg is the solubility of the gas in the solution
phase
• Pg is the partial pressure of the gas over the
solution
• k is a proportionality constant
Concentration Stuff
• Mass %=Mass of component is solution/ total
mass solution
• PPM: Parts per million (10^6)
• Molarity=Moles solute / liters solution
• Molality= moles of solute / kg solvent
Colligative Properties
• Boiling-point elevation:
• Freezing-point depression:
• Osmosis: pi=(n/v)RT=MRT
– The net movement of solvent is always toward the
solution with the higher solute concentration
Raoult’s Law
• Pvapor=XP°vapor
– X is the mole fraction of a solvent in solution
– P°vapor is the vapor pressure of the pure solvent
– Pvapor is the partial pressure of a solvent over a
solution
– = is the equal sign
Vocab
• Volatile: Liquids that evaporate readily
• Miscible: Pairs of liquid that mix in all
proportions
• Journalism: No science involved!