13.S: Properties of Solutions (Summary)
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- 70553
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
- \(\Delta H_{\mathit{soln}}=\Delta H_1+\Delta H_2+\Delta H_3\)
- 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:
- \(C_g = kP_g\)
- Cg solubility of gas in solution phase (usually expressed as molarity), Pg 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:
\[\displaystyle\textit{mass% of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 100 \nonumber \]
- very dilute solutions expressed in parts per million (ppm)
\[\displaystyle\textit{ppm of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 10^6 \nonumber \]
- 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
\[\displaystyle\textit{mole fraction of component} = \frac{\textit{moles of component}}{\textit{total moles of all components}} \nonumber \]
- 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: \(P_A = X_AP_{A}^°\)
- PA = vapor pressure of solution, XA = mole fraction of solvent, PA° = 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
- \(\Delta T_b=K_b m\)
- Kb = 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:
- \(\Delta T_f=K_f m\)
- Kf = 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