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 H 3 has to be comparable to H 1 + H 2 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 O 2 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 (10 6 ) grams of solution or 1mg solute per kg solution 1ppm = 1mg solute/L solution 1 ppb = 1g of solute/10 9 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: \(P_A = X_AP_{A}^°\) 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