This document discusses chemical mixtures and solutions. It begins by defining key terms like solvent, solute, and miscibility. Intermolecular forces like Coulombic forces, van der Waals forces, and hydrogen bonding determine whether substances will mix. The entropy and enthalpy of the solution process depend on the intermolecular forces between solute-solute, solvent-solvent, and solute-solvent particles. Temperature and pressure can impact gas solubility based on Henry's Law. Concentrated solutions are described using terms like molarity, molality, mole fraction and mass percent. Colligative properties like vapor pressure, freezing point, and boiling point depend only on the number of solute particles regardless of type

1. 1
Chemical Mixtures
• Many materials exist as homogeneous mixtures
– air (g), lake water (l), gasoline (l), stainless steel (s)
• Mixtures include suspensions, colloids, and
solutions
• Solution - homogeneous mixture of of two or more
substances; often comprised of solvent and solute, e.g., when
gases or solids are dissolved in a liquid
• Solvent - component with same phase as solution; substance
present in excess in liquid-liquid mixtures
• Solute - minor component mixed with solvent
• Degree of mixing of fluids
– Miscible - mix or dissolve in all proportions
– Immiscible - form layers upon mixing
• Solubility - maximal amount of solute accommodated
by solvent at specified temperature

2. 2
Molecular Forces
• All condensed matter has intermolecular attractive
forces (IAFs) between particles (formula units, ions,
molecules)
• Predominant attractive forces
• Coulombic - attraction of oppositely charged
particles
• van der Waals forces: a. dispersion (London);
b. dipole-dipole; c. ion-dipole
• Hydrogen-bonding
• Solutions form when it is energetically favorable
• Solutions form when substances with similar IAFs are
mixed: “like dissolves like,” compounds of like polarity
are most compatible

3. 3
Lattice and Hydration Energies
• Lattice Energy – energy of attraction between
counter ions in an ionic solid (Coulombic forces)
– Lattice bonds must be broken to separate ions in
solid
– IAF must be broken to separate solvent (water)
molecules; this involves breaking hydrogen bond
network between water molecules
• Hydration Energy – New IAF formed between
ions and water molecules; involves forming ion-
dipole bonds
– Water orients around released ions according to
polarity

4. 4
Energy Changes
• Heat of Solution (Hsoln); Heat flow in dissolution
solute-solute interactions (+H1 to break IFAs)
solvent-solvent interactions (+H2 to break IFAs)
solute-solvent interactions (-H3 to form IFAs)
Hsoln = H1 + H2 + H3
The sign of Hsoln (+ or -) depends on the magnitudes
of H’s 1-3
• Entropy of Solution (Ssoln); Randomness in dissolution
When substances are mixed, Ssoln is + because disorder
is increased (or order is decreased)

5. 5
Solution Process
• Factors which promote dissolution are:
– increase in randomness of system (entropy)
– release of heat upon dissolution (enthalpy), IAFs
• this occurs when solvent-solute IAF bonds are
very strong
• Dissolving an ionic solid in water
– water molecules collide with exterior ions of the solid
lattice
– ions break free from lattice and are hydrated by
water molecules, stabilizing the ion in solution
– extent of dissolution depends on lattice vs.
hydration energies

6. 6
Predicting Solubility
• Substances that are compatible have
similar IAFs (i.e., like dissolves like)
• Nonpolar substances (e.g., hydrocarbons)
have dispersion force IAFs
• Polar substances have dipole driven IAFs
(e.g., methanol)
• Substances with similar IAFs mix to the
greatest extent
– Ionic compounds are most soluble in polar
solvents

7. 7
Solution Nomenclature
• Solution Formation: Dissolving a solid in a liquid
• undersaturated solution (below solubility limit)
NaCl(s) Na+(aq) + Cl-(aq) ionic
C6H12O6(s)  C6H12O6(aq) molecular
• saturated solution (at solubility limit)
NaCl(s)  Na+(aq) + Cl-(aq) Dynamic Equilibrium
C6H12O6(s)  C6H12O6(aq)  =
reversible
rate of solution = rate of crystallization; heterogeneous mixture
forms
s = constant @ T with introduction of more solute
• supersaturated solution - solution which contains
more than saturated concentration of solute

8. 8
Effect of P on Solutions
• P changes do not greatly effect solubility of liquids
and solids significantly aside from extreme P changes
(vacuum or very high pressure)
• Gases are compressible fluids and therefore
solubilities in liquids are P dependent
• Solubility of gas in liquid  partial pressure of same
gas (only) above solution
• Quantitatively defined in terms of Henry’s Law:
Cg = kPg
Cg = solubility of gas (mol/L); Pg = partial pressure of
gas (atm); k = Henry’s Law Constant (mol/L-atm)

9. 9
Effects of T on Solutions
• Solubility of the solute is temperature dependent
• Gaseous solutes in liquid solvents
 T leads to solubility (& vice versa)
A(g) A(l)  A(aq) H < 0
• Solid solutes in liquid solvents
– most salts:  T leads to solubility (& vice versa)
A(s) A(l)  A(aq) H > 0
– some ionic compounds:  T leads to solubility (&
vice versa)
A(s) A(l)  A(aq) H <0
• Liquid solute behavior in liquid solvents is unpredictable

10. 10
Concentration Terms
solvent
of
kg
solute
mol
(m)
Molality =
0.5 mol of solute in 1 kg of solvent = 0.5 molal solution
molality is not temperature dependent
solution
of
mol
solute
of
mol
(X)
fraction
Mole =
1 mol of solute (A) in 3 mol of solvent (B) = 0.25
Mole Fraction Convention:
A = solute and B = solvent

11. 11
Concentration Terms
solution
L
solute
mol
(M)
Molarity =
0.5 mol of solute in 0.5 L of sol’n = 1 molar solution
molarity is temperature dependent
100
solution
of
mass
solute
of
mass
(%)
percent
Mass 
=
1 g of solute in 100 g of solution = 1% solution
Remember that solution mass or volume =
mass/volume of solute + mass/volume of solvent

12. 12
Colligative Properties of Solutions
• Important solution properties which show colligative
behavior
– vapor pressure, Pv
– freezing point, Tf
– boiling point, Tb
– osmotic pressure, 
• Colligative properties: solution properties which
depend only on the relative number of solute to solvent
particles (solute concentration) and not on the chemical
properties of the solute
• Differences between solvent and solution illustrated in
phase diagram (P versus T plot)

13. 13
Colligative Mathematical
Property Relation
1. vapor pressure
(Raoult’s Law)
2. freezing point depression
3. boiling point elevation
4. osmosis
o
A
A
A P
X
P =
m
K
T f
f =

m
K
T b
b =

MRT
=


14. 14
Vapor Pressure, Pv, and Vapor
Pressure Lowering, Pv
• Vapor pressure: partial pressure of vapor above
liquid
• Solution vapor pressure
– Examples include volatile solvent with non-volatile
solute and volatile solute and volatile solvent
– In example of non-volatile solute, vapor pressure
of solution proportional to concentration of solvent
Pv,A  [solvent A] = XAPo
A
As solvent concentration decreases, Pv decreases
– In example of volatile solute (B) and solvent (A)
PT = PA + PB = XAPo
A + XBPo
B

15. 15
14
Semipermeable membrane
  
 
  
 
  
 
  
 
= solute molecule
 = solvent molecule
1 2
•Osmosis = solvent flow across membrane
•Net movement of solvent is from 2 to 1; pressure increases
in 1
•PA,1 < P A,2, osmosis follows solvent vapor pressure
•Osmotic Pressure = pressure applied to just stop osmosis
2 Compartment Chamber