I. Solutions - HOMOGENEOUS MIXTURES
A. Homogeneous Mixtures
1. The particles of each substance in the mixture are evenly dispersed throughout the mixture.
2. The solution is stable and the dissolved material in the solution will not spontaneously separate
from the mixture.
3. The dissolved particles, which can be ions (Na+, Cl-) or molecules (C12H22O11) are individually in
contact with the water molecules, not clustered together as in non-examples of mixtures.
a. alloys – solutions of metal atoms (brass, bronze, pewter, steel, 18 karat gold)
b. air - gaseous solutions (solvent = nitrogen @79%, solutes = O2, Ar, CO2)
c. brine – salt water
d. tinctures – solvent is alcohol
B. NON - EXAMPLES are Heterogeneous mixtures.
1. Suspension - a mixture that appears uniform while being stirred, but separates into
different phases when agitation ceases.
a) chocolate syrup in chocolate milk, muddy water
2. Colloids – a mixture in which small particles are suspended throughout the solvent
a) the particles are small but too large to dissolve
b) dispersed phase - the particles that do not dissolve
c) dispersing medium – analogous to the solvent
d) Ex: milk with small fat globules and small lumps of the protein casein dispersed in the clear
C. Describing Solution – in the form of a homogeneous liquid mixture as opposed to gaseous (air) & solid
1. Solvent - the material dissolving the solute in a solution
a) Water - the most common solvent – aqueous solutions
b) Alcohol – solutions are called tinctures (useful when solute is insoluble in water)
2. Solute - the material dissolved in a solution.
a) may be solid, liquid or gas solutes
1. Solid solute in a liquid solvent - most common
a) ex. Salts dissolve and their ions are attracted to the dipoles of the solvent
(remember water is polar) making a stable, homogeneous mixture.
2. Liquid solute in a liquid solvent
a) if the two liquids mix together (homogeneously) they are miscible.
1) Miscible: liquids or gases that will dissolve in each other.
2) Remember: “Likes dissolve Likes”
b) If the two liquids do not mix together (homogeneously) it is immiscible.
1) Immiscible: liquids or gases that will not dissolve in each other.
2) Polar and Nonpolar liquids
3. Gas solute in a liquid solvent
a) ex. carbonated water & ammonia cleaning solutions.
D. Forming a solution
1. The extent to which substances form solutions is controlled by the relative strengths
of their IMFA.
2. If a solution is to form:
a. Energy must break the forces of attraction between the solvent.
b. Energy must break the forces of attraction between the solute.
c. Energy must be released as the particles are brought together. (called
3. The H solution or Enthalpy of Solution = Net change in energy in forming a solution.
a. If H solution = + value, then the energy needed to overcome forces of
attraction in the crystal was greater than the energy released in hydrating the
1) the solution cools down as the solute dissolves.
2) Endothermic change
b. If H solution = - value, then the energy needed to overcome forces of attraction
in the crystal was less than the energy released in hydrating the solute.
1) the solution heats up as the solute dissolves.
2) Exothermic change
c. Sample Data
Solid Lattice Energy Hhydration Hsolution Solubility in H2O
(kJ/mol) (kJ/mol) (kJ/mol) (g/100 mL H2O)
AgCl 912 -851 +61 0.00089 @ 10oC
Cation size increases NaCl 786 -760 +26 36.7 @ 0oC
down alkali metal family. LiF 1037 -1005 +32 0.3 @ 18oC
Thus, shielding is greater
and it is easier to separate
KF 821 -819 +2 92.3 @ 18oC
the ions from each other. RbF 789 -792 -3 130.6 @ 18oC
H hydration is Only the last Solubility increases as
greater for the larger solid releases H solution decreases
ions so less energy is heat as it (more exothermic)
released upon dissolves
II. Concentration: the ratio of solute to solvent in a solution. (A qualitative term)
A. Many different ways to Quantitatively express concentration:
Concentration Name Abbreviation or Symbol Units Areas of Application
Solubility g/100 g g solute Determining solubility’s
100 g solvent and in medical products
Mass Percent % or % w/w g solute x 100 In biological research
Parts per million ppm g solute EPA, small concentrations as
1,000,000 g solution in pollutants or contaminants
Parts per billion ppb g solute EPA, small concentrations as
1,000,000,000 g solution in pollutants or contaminants
Parts per trillion ppt g solute Isotopes that are used as
1,000,000,000,000 g soln tracers in medicine
Concentration Name Abbreviation or Symbol Units Areas of Application
Molality m mol solute Colligative properties i.e.,
kg solvent fp depression & bp elevation
Mole Fraction X mol solute In solution
mol solute + mol solvent thermodynamics
Volume percent % V/V mL solute x 100 Liquid-liquid mixtures
Mass-volume percent % w/v g solute x 100 In many commercial
mL solution products
B. MOLARITY, M = mol/L mole solute/L solution
1. Four different labels: Molarity, Molar, M or mol/L
2. Molarity is defined in terms of the volume of the entire solution NOT just the volume of solvent
a) Therefore you must account for the vol. of the solute and the vol. of the solvent when combined
b) Sometimes upon adding solute to a solvent the volume increases or decreases. The final volume
is generally 1 L.
C. Properties of Solutions
a. g solvent g solution, because mass of solution is mass of solute + solvent
b. vol. solvent vol. solution, because the number and type of interactions between
particles in solution the sum of the number and type of
interactions in the solute and solvent
c. T means a M, because as a solution is heated, the intermolecular distance between
molecules changes. Thus, molarity changes with T
d. T means m, because mass is conserved and as the solution is heated the # kg of water
HW : Find M, m, mass percent and X of a 10oC, 1 L KCl solution which has a solubility of
30.0 g KCl /100g water and a density of 1.01 g/mL.
Answers: m = mole solute m = 30 g KCl x 1 mol KCl x 1000 g = 4.0 m
Kg of solvent 100 g H2O 74.55 g KCl 1 kg
M = mol solute = 30.0 g KCl x 1mol KCl / 74.55 g KCl = 0.402 mol KCl = 3.12 M
L solution = 100 g H2O + 30 g KCl or 130 g soln x 1 L/ 1000 mL = 0.129 L soln
1 mL KCl 1L
Mass % = 30 g KCl x 100 = 23%
130 g soln
X= mole solute = 0.405 mol KCl
Mol solute + mol solvent 0.405 mol KCl + 100 g H2O/18.02 g/mol H2O
= 0.405 mol KCl = 0.067
0.405 mol KCl + 5.555 mol H2O
C. Separating Mixtures
1. Decanting – separate top lay b/c of less density - separates a heterogeneous mixtures
2. Centrifugation – spinning at high speeds separates based on density - separates a heterogeneous mixtures
3. Filtration – based on varying degrees of solubility - separates a heterogeneous mixtures
4. Crystallization – gradual evaporation of solvent from solute - separates a homogeneous mixtures
5. Chromatography – separating liquids or solids based on varying degrees of attraction to a
6. medium (IMFA) - separates a homogeneous mixtures
a) paper chromatography – capillary action – stronger IMFA to paper moves slower
b) gas chromatography
7. Distillation – evaporation and then condensation - separates a homogeneous mixtures
III. Solubility – A reference to the maximum solute that can dissolve in a solvent
1. Soluble - a solute that completely dissolves in a solvent
2. Insoluble - a solute that cannot dissolve in a solvent.
3. The differences between soluble, partly soluble and insoluble are vague.
a. When more precision is needed, the exact amount of a solute that can dissolve in a solvent
can be expressed numerically (at some temp.)
b. Solubility - the maximum amount of a chemical that will dissolve in a given amount of
solvent at a specified temperature while the solution is in contact with some
1) Units g/100mL water or g/100g water 1 mL = 1 g for water
2) Refer to solubility curves WS
3) Temperature affects solubility
a) Increased temp usually increases the solubility (for solid solutes)
b) BUT some solutes (mostly gases) show a decrease in solubility as the
Ex. SO2, Li2SO4
4) Pressure affects solubility of gases
a) Henry’s Law – The solubility of a gas in a liquid is directly proportional to
the partial pressure of that gas on the surface of the liquid.
1) The greater the partial pressure in a bottle of soda above the liquid, the
more CO2 that will dissolve in the liquid.
2) As the bottle is opened, partial pressure of CO2 quickly drops to that of
atmospheric pressure and thus less gas can be dissolved.
3) Practical application of this occurs with divers – the bends
4) Formula Sg = KH . Pg Where Sg = solubility of a gas
KH = Henry’s constant which is
Characteristic of the solute
And solvent and changes w/
Pg = pressure of the gas
a) The magnitude of KH depends on the strength of the IMFA
between the gas and the solute.
b) Ex. O2 vs. N2
1) Oxygen gas is larger thus greater electron-electron
interaction (greater polarizable nature). So with the
stronger dipole-induced forces O2 is more soluble than N2
4. The Dissolving Process is a reversible process.
a. While some ions are dissolving, others recrystallize
b. As more ions dissolve the rate of recrystallization also increases.
c. At the maximum solubility:
the rate of dissolving = rate of recrystallizing
d. An Unsaturated solution is one in which there is less solute than the specified solubility amount
at a given temperature.
1) more solute will dissolve when added
e. A Saturated solution is one in which the specified amount of solute that a solution can hold at a
given temperature has been met.
1) equilibrium system – rate of dissolving = rate crystallization (undissolving) @ constant T
2) No more solute will dissolve at that temperature. Additional solute sinks to the bottom.
f. A Supersaturated solution contains more than the standard amount of solute specified by the
solubility at a given temperature
1) No more solute dissolves, and as one crystal is added to the solution, all of the excess solute
that had dissolved precipitates out of solution.
2) These solutions do NOT break the rules b/c they are NOT in contact w/ undissolved solute.
3) Hot packs: Supersaturated solutions like (NaC2H3O2) when activated or disturbed have their
dissolved solute precipitate out of solution releasing great amounts of energy as the mixture
approaches a more stable state. (To reuse, reheat to form dissolve solute, re-supersaturate)
5. Some substances (ionic compounds) do not dissolve: Refer to Rules of Solubility
6. Affecting Solubility
a. Increasing the temperature:
1) increases the solubility for nearly all solids,
a) exceptions Li2SO4 and Ce2(SO4)3
2) decreases the solubility of all gases.
b. Le Chateliers’ Principle – states that a change in any of the factors determining the
equilibrium of a system causes a system to counteract the affect of the change.
1) Ex. To increase Temperature adds heat to a system and affects equilibrium.
a) Reaction: NaC2H3O2 (s) + H2O(l) + Heat NaC2H3O2 (aq)
increase in heat shifts reaction to right forming
a supersaturated solution
This supersaturated solution is very unstable so any disturbance
will shift the reaction in the reverse direction.
b) Reaction: Gas (solvent) + Liquid (solvent) Saturated solution + heat
if heat is increases the rxn shifts
to the left and gas escapes
c. Colligative Properties – properties of a solvent that are affected by the amount of solute that
is present rather than the kind or chemical properties of the solute.
a. The solute interferes with the normal actions of the solvent’s molecules.
b. The [solute] or concentration of the solute determines the property.
c. In other words, 0.5 g of NaCl will affect 1 liter of water the same as 0.5 g of KI.
2. Vapor Pressure – is a colligative property
a. The addition of solute to solvent lowers the Vapor Pressure
b. Fewer solvent molecules can get to the surface in order to evaporate so the vapor
pressure drops with increased solute.
c. Raoult’s Law: The vapor pressure of a solution is proportional to the fraction
solvent molecules at
the surface and will equal the mole fraction of the solvent in the solution.
Psolvent = Xsolvent . Posolvent
Vapor Pressure of the pure solvent
Mole fraction of solvent (convert g mol solvent / g mole solute + solvent)
Vapor pressure of solvent in solution
3. Boiling Point elevation – Boiling point of a solvent increases with increased [solute]
a. Results from a lower vapor pressure
b. Remember a liquid boils when its vapor pressure = atmospheric pressure
c. If the vapor pressure is lower due to solutes, it takes a higher temperature to get the
liquid molecules to reach the atmospheric pressure and escape as a gas.
Tbp = Kbp . m . i where i is the # of particles /molecule (ex. NaCl i is 2 Na+ & Cl-)
(ex. Sugar, C12H22O11 i is 1 molecule dissolves as 1 unit )
molality of solute = # mole of solute / 1 kg of solvent
boiling point elevation constant (for water = 0.5121 oC/molal) This value
is unique for each solvent.
Change in temperature for boiling which is added to the original bp to get new bp.
2. Freezing point depression – the freezing point of a solvent decreases with increased
Tfp = Kfp . m . i where i is the # of particles /molecule (ex. BaCl2 i is 3 Ba2+ & 2Cl-)
molality of solute = # mole of solute / 1 kg of solvent
freezing point elevation constant (for water = -1.86 oC/molal) This value
is also unique for each solvent.
Change in temperature for freezing which is subtracted from the original fp to get new fp.
3. Osmotic Pressure – The pressure created as pure solvent molecules passes through a
semi-permeable membrane into a solution of high concentration
until counteracting pressure (called osmotic pressure) stops the
a. Osmosis occurs until equilibrium of concentrations is reached.
b. Symbol for osmotic pressure is the upper case Greek letter for pi =
= c . R . T where T is temperature in Kelvin
the Gas constant
concentration of the solution
c. Ex. Egg (with a semi-permeable shell) in water – osmotic pressure forces
water into the egg since there is a higher concentration of solutes inside the
egg. Thus the egg swells
d. Ex. Egg – in corn syrup – Higher concentration outside egg, osmotic
pressure forces solvent out of egg into corn syrup – egg’s volume decreases.
7. Rules for Solubility of Ionic Compounds:
The following rules must be applied in the given order. The smaller number takes precedence in case
of a conflict.
1. All Group I elements and Ammonium salts are soluble.
2. Nitrates, Acetates, Chlorates and Perchlorates are soluble.
3. Silver, Lead and Mercury I salts are insoluble.
4. Fluorides, Chlorides, Bromides and Iodides are soluble.
5. Carbonates, Silicates, Phosphates, Oxalates & Chromates are insoluble.
6. Sulfides, Oxides and Hydroxides are insoluble (Except: Ca(OH)2 (aq), Sr(OH)2 (aq) & Ba(OH)2 (aq)
which are soluble.)
7. Sulfates are soluble (Except CaSO4(s), SrSO4(s) & BaSO4(s) which are insoluble.)
8. Predicting Solubility
a. Likes Dissolve Likes
1. Polar substances will dissolve in other polar substances but not in nonpolar substances
a) Solutes that form hydrogen bonds with the solvent have high solubility.
2. Nonpolar substances will dissolve only in other nonpolar substances.
b. Molecular polarity Review
1. Molecules are polar if:
a) they have an uneven distribution of atoms (ASYMMETRICAL)
ex. angular/bent molecules ie group 16 O, S
pyrimidal/tripod molecules ie group 15 N, P
b) they have different atoms attached to a central atom
c) if they have large difference in the electronegativities of the bonded atoms.
1) polar covalent molecules have difference b/w 0.5 and 2.1
2. Molecules are nonpolar if:
a) they have and even distribution of atoms - SYMMETRICAL
ex. F-F, BH3, CH4
b) they have all the same atoms attached to a central atom.
c) they have very small differences b/w the electronegativities of the bonded atoms.
1) differences b/w 0.0 and 0.4
9. Using the Principles of Solubility
a. Dry Cleaners use these principles to clean stains.
1. C2Cl4 - tetrachloroethylene
a. Good, non-polar solvent that is not flammable and is very volatile, thus recycles
(up to 200 times)
b. Bad, carcinogen, Health officials recommend it not be inhaled
c. Polar stains are pretreated with a polar solvent and then allowed to dry. Upon
drying they are then treated with the nonpolar dry-cleaning agent.
b. Vitamins can be water soluble (polar) or fat soluble (nonpolar)
1. Water soluble vitamins are regulated by the kidneys.
ex. B complex and C
2. Fat soluble vitamins link with special proteins to travel in the blood.
ex. A, D, E, K
IV. Electrical Conductivity of Solutions
A. Solutions with Free Flowing Ions are capable of conducting electricity
1. Electrolyte – a substance that, when dissolved in a solvent, increases the solvents
a. Strong electrolyte – dissociate completely in solution and provide high conductivity
1) strong acids – those molecules that dissociate completely into H+ ions
2) salts -all ionic compounds to some degree have some ions that dissolve in
solution and thus have free flowing ions to conduct
b. Weak electrolyte - dissociate only partly in solution and provide weak conductivity
1) ex. weak acids – only partially dissociate into H+ ions
2. Nonelectrolytes - a substance that, when dissolved in a solvent, does not enhance the
solvents ability to conduct electricity.
a. polar molecules that dissolve in water without dissociating into ions
B. Tap Water conducts electricity b/c water contains dissolved minerals and ions from water treatment
1. It is a very, very, very, very weak electrolyte.
V. Overcoming Insolubility
A. Imagine: Oil (nonpolar) and Vinegar (polar) - Immiscible
1. Less dense oil rests on top of the vinegar.
2. BUT what if we shake the container? Agitation keeps the suspension from settling
a) Millions of tiny, microscopic droplets of vinegar are surrounded by oil.
b) This mixture is called an EMULSION
- small, insoluble particles called colloids in the form of droplets of one liquid suspended in
3. NOW, there are commercial salad dressings made of oil and vinegar that are permanent emulsions
which do not separate back into layers after mixing.
a. Lecithin - a family of several similar molecules that stabilize emulsions.
1) obtained from soybeans
2) By itself, lecithin dissolve in oils, not in water (polar substance)
3) BUT in the presence of both oil & water, each molecule of lecithin dissolve partly in
the oil and partly in the water, thus linking the two immiscible liquids.
4) it is an example of an EMULSIFYING AGENT.
b. Soap – an Emulsifying agent which is a sodium or potassium salt of a long-chain fatty acid
1) perspiration is a mixture of water & body oils.
a) water cools the skin, oils keep it soft.
2) Water alone cannot remove the build up of oils that embed dirt & bacteria.
3) Soaps form micelles that emulsify oil and water colloid dispersions to help clean the
skin. (Micelles are a combination of polar/nonpolar particles)
a) nonpolar end of soap dissolves the oils
b) polar end of soap dissolves in the water.
c) Surfactants – a class of salts, valued for their cleansing properties whose
anions possess a negatively charged “head” and a long nonpolar “tail”
1) a negatively charged end is “hydrophilic”, or water loving and is thus,
soluble in water
2) the other end which is nonpolar and “hydrophobic” is insoluble in H2O
but attracts the nonpolar oils or matter.
3) Thus, a surfactant, or surface active agent, acts to stabilize an
emulsion by acting at the surface b/w two immiscible substances.
4) Essentially, they coat the surfaces of the dispersed phase particles thus
preventing them from adhering to one another
4) Hard water (water containing high [ ] of Ca2+ ions, Fe2+ions & Mg2+ ions ) destroy the
emulsifying ability of soap.
a) these ions cling to the polar ends of the soap making the soap less effective.
5) Detergents - soap substitutes that were discovered in the 1930's which work well in
hard water as emulsifying agents.
a. Add PO4 –3 to solution which bond to Ca2+ ions, Fe2+ions & Mg2+ which ppt
out of solution allowing soaps to clean better)