13.8: Properties of Solutions (Summary)

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These is a summary of key concepts of the chapter in the Textmap created for "Chemistry: The Central Science" by Brown et al.

13.1: The Solution Process

interaction between solute and solvent molecules
hydration – solvation when solvent is water

13.1.1 Energy Changes and Solution Formation

overall enthalpy change in formation of a solution
- \(\Delta H_{\mathit{soln}}=\Delta H_1+\Delta H_2+\Delta H_3\)
  - \(\Delta H_1=\) separation of solute molecules
  - \(\Delta H_1=\) separation of solvent molecules
  - \(\Delta H_3=\) formation of solute-solvent interactions
separation of solute particles is endothermic
separation of solvent is endothermic
third is exothermic
formation of solution can be either exothermic or endothermic
exothermic processes are spontaneous
solution will not form if enthalpy is too endothermic
H₃ has to be comparable to H₁ + H₂
- Ionic substances cannot dissolve in nonpolar liquids
- Polar liquids do not form solutions with nonpolar liquids

13.1.2 Solution Formation, Spontaneity, and Disorder

two nonpolar substances dissolve in one another
attractive forces = London dispersion forces
two factors in processes that are spontaneous: energy and disorder
processes in which the energy content of the system decreases tend to occur spontaneously
- exothermic
processes in which the disorder of the system increases tend to occur spontaneously
solutions will form unless solute-solute or solvent-solvent interactions too strong relative to solute-solvent interactions

13.1.3 Solution Formation and Chemical Reactions

distinguish between physical process of solution formation from chemical process that leads to a solution

13.2: Saturated Solutions and Solubility

crystallization – reverse process of solution
dynamic equilibrium – when equilibrium exists between process of solution and crystallization
solute said to be saturated
solubility – amount of solute needed to saturate a solution
unsaturated – when there isn’t enough solute to saturate a solution
supersaturated – when there is more solute than needed to saturate a solution
for most salts crystallization of excess solute is exothermic

13.3: Factors Affecting Solubility

Solute-Solvent Interactions

solubility increases with increasing molar mass
London dispersion forces increase with increasing size and mass of gas molecules
Miscible – pairs of liquids that mix in all proportions
Immiscible – opposite of miscible
Hydrogen-bonding interactions between solute and solvent leads to high solubility
Substances with similar intermolecular attractive forces tend to be soluble in one another
"like dissolves like"

Pressure Effects

solubility of a gas in any solvent increases as pressure of gas over solvent increases
relationship between pressure and solubility: Henry’s Law:
- \(C_g = kP_g\)
- C_g solubility of gas in solution phase (usually expressed as molarity), P_g partial pressure of gas over solution, k is proportionality constant (Henry’s Law constant)
- Henry’s law constant different fore each solute-solvent pair, and temperature

Temperature Effects

solubility of most solid solutes in water increases as temperature of solution increase
solubility of gases in water decreases with increasing temperature
decreases solubility of O₂ in water as temperature increases in one the effects of thermal pollution

13.4: Ways of Expressing Concentration

dilute and concentrated used to describe solution qualitatively
mass percentage of component in solution:

\[\displaystyle\textit{mass% of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 100\]

very dilute solutions expressed in parts per million (ppm)

\[\displaystyle\textit{ppm of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 10^6\]

1ppm = 1g solute for each (10⁶) grams of solution or 1mg solute per kg solution
1ppm = 1mg solute/L solution
1 ppb = 1g of solute/10⁹ grams of solution or 1 m g solute/ L of solution

13.4.1 Mole Fraction, Molarity, and Molality

\[\displaystyle\textit{mole fraction of component} = \frac{\textit{moles of component}}{\textit{total moles of all components}}\]

sum of mole fractions of all components of solution must equal one
\(\displaystyle\textit{molarity} = \frac{\textit{moles solute}}{\textit{liters soln}}\)
\(\displaystyle\textit{molality} = \frac{\textit{moles solute}}{\textit{kilograms of solvent}}\)
molality goes not vary with temperature
molarity changes with temperature because of expansion and contraction of solution

13.5: Colligative Properties

Colligative properties are physical properties that depend on quantity

Lowering the Vapor Pressure

vapor pressure over pure solvent higher than that over solution
vapor pressure needed to obtain equilibrium of pure solvent higher than that of solution

Raoult’s Law

Raoult’s law: \(P_A = X_AP_{A}^°\)
- P_A = vapor pressure of solution, X_A = mole fraction of solvent, P_A° = vapor pressure of the pure solvent
- Ideal solution – solution that obeys Raoult’s law
- Solute concentration is low, solute and solvent have similar molecular sizes and similar types of intermolecular attractions

Boiling-Point Elevation

normal boiling point of pure liquid is the temperature at which pressure is 1 atm
addition of a nonvolatile solute lowers vapor pressure of solution
\(\Delta T_b=K_b m\)
K_b = molal boiling-point-elevation constant
- Depends only on solvent
- boiling point elevation proportional to number of solute particles present in given quantity of solution

Freezing-Point Depression

freezing point of solution is temperature at which the first crystals of pure solvent form in equilibrium
freezing point of solution lower than pure liquid
freezing point directly proportional to the molality of the solute:
- \(\Delta T_f=K_f m\)
- K_f = molal freezing-point-depression constant

Osmosis

semipermeable – membranes that allow passage of some molecules and not others
osmosis – the net movement of solvent molecules from the less concentrated solution to the more concentrated solution
net movement of solvent always toward the solution with the higher solute concentration
osmotic pressure – pressure needed to prevent osmosis, p
- \(\pi = \left(\frac{n}{V}\right)RT=MRT\)
- M = molarity of solution
if solutions identical osmosis will not occur and said to be isotonic
if one solution lower osmotic pressure = hypotonic, the solution that has higher osmotic pressure = hypertonic
crenation = when cells shrivel up from the loss of water
hemolysis = when cells rupture due to to much water

Determination of Molar Mass

colligative properties can be used to find molar mass

13.6: Colloids

Colloidal dispersions (colloids) are intermediate types of dispersions or suspensions

intermediate solutions between solutions and heterogeneous mixtures
colloids can be gases, liquids, or solids
colloid particles have size between 10 - 2000Å
tyndall effect – scattering of light by colloids

Hydrophilic and Hydrophobic Colloids

hydrophilic – colloids in which the dispersion medium is water
hydrophobic – colloids not dispersed in water
hydrophobic colloids have to be stabilized before being put in water
- natural lack of affinity for water causes separation
- can be stabilized by the adsorption of ions on the surface
- adsorped ions interact with water
- can also be stabilized by presence of other hydrophilic groups on surface

Removal of Colloidal Particles

coagulation – enlarging colloidal particles
- heating or adding an electrolyte to mixture
- heating increases number of collisions and particles stick together increasing their size
- electrolytes causes neutralization of the surface charges of the particles which remove the electrostatic repulsion
- dialysis – use of semipermeable membranes to filter out colloidal particles