Solutions
Under the supervision of Dr.
Mahjoub Hashem, PhD in Chemistry.
Ibrahim Yusuf Ibrahim
1422001
3/D
What’s a solution?
Solution
1. A solution is a type of homogeneous mixture with uniform
composition.
2. Its particles are molecule-sized, making it appear consistent even
under a microscope.
3. Solutions form when one substance (solute) dissolves completely in
another (solvent).
4. Examples include saltwater, where salt dissolves evenly in water.
5. Unlike suspensions, solutions do not separate over time and remain
stable.
Understanding Mixtures:
Solutions and Suspensions
Solutions and Suspensions
1. Mixtures can be either heterogeneous (visible
particles) or homogeneous (uniform, molecule-
sized particles).
2. A homogeneous mixture appears uniform and is
also known as a solution.
3. Suspensions are temporary heterogeneous
mixtures where particles settle over time.
4. In a suspension, particles may remain mixed
temporarily but eventually separate and settle.
5. An example of a suspension is clay in water,
where clay settles at the bottom over time.
The Energies of Solution Formation
The Energies of Solution
Formation
1. Solubility follows the rule "like dissolves like": polar
solvents dissolve polar/ionic solutes, and nonpolar
solvents dissolve nonpolar solutes.
2. Solution formation involves three steps: separating
solute particles, expanding the solvent, and allowing
solute-solvent interactions.
3. The enthalpy of solution (ΔHsoln) is the sum of
energy changes in these steps, which can be
endothermic or exothermic.
4. NaCl dissolves in water despite a small positive
ΔHsoln because the mixed state is more probable and
energetically favorable.
5. Nonpolar solutes (e.g., oil) do not dissolve in polar
solvents (e.g., water) due to large positive ΔHsoln
from weak solute-solvent interactions.
Describing Solution
Composition
Solution composition can be described
qualitatively (dilute vs. concentrated) or
quantitatively using precise measures like
molarity (M).
Mass percent is the percent by mass of solute
in the solution, while mole fraction (χ) is the
ratio of moles of a component to total moles
in the solution.
Molality (m) is the number of moles of solute
per kilogram of solvent, useful for
temperature-dependent calculations.
Normality (N) measures equivalents per liter,
where equivalents depend on the reaction
(e.g., acid-base or redox).
Factors Affecting Solubility
Pressure Effects
1. Pressure significantly increases gas
solubility (e.g., CO₂ in sodas) but has little
effect on solids/liquid
2. Higher pressure forces more gas into
solution until a new equilibrium is reached
3. Henry’s Law – The concentration (C) of
dissolved gas is directly proportional to its
partial pressure (P): C = kP.
4. Carbonated drinks are bottled under high
CO₂ pressure; opening the can releases gas due
to lower pressure.
5. Henry’s law applies best to gases that don’t
react/dissociate in the solvent (e.g., O₂ in
water, but not HCl).
Temperature Effects
1. In solids solubility increases as
temperature increases
2. In gases solubility decreases as
temperature increases.
3. The temperature dependence of
solubility is complex and must be
determined experimentally
4. Warmer water holds less oxygen, creating
stratified layers that block oxygen
absorption in deep lakes, endangering marine
life.
5. Heating water expels CO₂, shifting
equilibrium to produce insoluble calcium
carbonate (CaCO₃), which forms damaging
deposits in pipes and boilers.
The Vapor Pressures of solutions
Vapor Pressure
1. Nonvolatile solutes alter solvent properties (e.g.,
lower freezing point in antifreeze, salt melting
ice).
2. Solutes decrease solvent vapor pressure by
reducing surface solvent molecule
concentration.
3. Raoult's Law: Solution vapor pressure (P) equals
solvent mole fraction (X) times pure solvent's
pressure (P°): P = XₛₒₗᵥP°ₛₒₗᵥ.
4. Pure solvent transfers to solution until vapor
pressures equalize (demonstrated in sulfuric
acid/water experiment).
5. Solute particles displace solvent molecules at
surface, proportionally reducing vaporization
tendency.
Raoult’s law
where Psoln is the observed
vapor pressure of the solution,
χsolvent is the mole fraction of
0
solvent, and P solvent solvent is
the vapor pressure of the pure
solvent.
Colligative Properties
Boiling-Point
Elevation
Boiling Point Definition - Normal boiling point
occurs when vapor pressure equals 1 atm.
Solute Effect - Nonvolatile solutes lower
solvent's vapor pressure, requiring higher
temperatures to boil.
Boiling Point Elevation - Solutions with
nonvolatile solutes boil at higher
temperatures than pure solvents.
Phase Diagram Shift - Solution liquid/vapor
equilibrium line shifts to higher temperatures
compared to pure solvent.
Concentration Dependence - The degree of
boiling point elevation depends on solute
concentration.
Freezing-Point
Depression
Freezing Point Depression - Solutes lower
solution's freezing point below pure solvent's
(e.g., saltwater freezes below 0°C).
Vapor Pressure Mechanism - Solute reduces
liquid's vapor pressure, requiring lower
temperature to match ice's vapor pressure.
Molecular Explanation - Solute disrupts
solvent's solid formation, slowing
crystallization until new equilibrium is
reached.
Practical Application - Salt (NaCl/CaCl₂) melts
ice by lowering water's freezing point
(ineffective at extremely cold temperatures).
Phase Diagram Impact - Solutes shift liquid
range: elevate boiling point AND depress
freezing point simultaneously.
Osmotic Pressure
Osmosis Definition - Solvent flows through a
semipermeable membrane from low to high
solute concentration until equilibrium is
reached.
Osmotic Pressure - The excess pressure
needed to stop osmosis equals the solution's
osmotic pressure (a colligative property).
Medical Applications - Dialysis uses
semipermeable membranes to remove waste
from blood (artificial kidneys).
Biological Significance - Isotonic solutions
maintain cell shape; hypertonic causes
crenation, hypotonic causes hemolysis.
Food Preservation - Salt/sugar create
hypertonic environments that kill bacteria by
dehydration (crenation).
Electrolyte
Solutions
Particle Concentration Dependence -
Colligative properties depend on total solute
particles.
Van't Hoff Factor (i) - Expected i = ions per
formula unit; actual i is often lower due to ion
pairing.
Ion Pairing Effect - Oppositely charged ions
temporarily pair, reducing effective particle
count (0.10 m NaCl shows i=1.87 not 2).
Dilution Matters - Ion pairing decreases in
dilute solutions (0.0010 m NaCl shows i≈1.97,
near ideal value).
Electrolyte Calculations - Colligative
equations for electrolytes include i: ΔT = i·Kₙ·m
(accounts for incomplete dissociation).
Colloids
Tyndall Effect - Light scattering by suspended
particles distinguishes colloids (visible beam)
from true solutions (invisible beam).
Colloid Definition - Stable suspensions of
particles (1-1000 nm) in a medium, classified by
dispersed/dispersing phases (e.g., solid in liquid).
Colloid Stability - Charged particle surfaces with
ionic layers create repulsive forces that prevent
aggregation and settling.
Coagulation Methods - Colloids can be destroyed
by heating (increases particle collisions) or
adding electrolytes (neutralizes charges).
Practical Examples - River delta formation (salt
coagulates clay) and electrostatic precipitators
(remove soot from industrial smoke).
