Solutions
Solutions Definition and Characteristics Homogeneous mixtures of two or more substances Appear to be pure substances Transparency Separation by filtration is not possible Uniform distribution of particles No settling or separation Uniform mixing is energetically favored by nature Composition may be varied
Most homogeneous mixtures we encounter can be described as a solution, i.e., a combination of a solute and a solvent.
Solute Solvent Substance oxygen nitrogen air water vapor air humid air carbon dioxide water soda water ammonia water household ammonia acetic acid water vinegar ethylene glycol water antifreeze salt water sea water sugar water syrup mercury silver dental amalgam zinc copper brass carbon iron steel copper silver sterling silver tin lead solder
Solvent and Solute in an Aqueous Solution
Solvent and Solute in an Aqueous Solution
Solution Formation When one substance (the solute) dissolves in another (solvent) it is said be soluble. When one substance does not dissolve in another it is said be insoluble. Solubility depends on two factors: tendency toward mixing (greater entropy) intermolecular forces
The Dissolving and Crystallization Processes Dissolving Crystallization The Dissolving Process Involves 1.Breakdown of the attractive forces between 2.Breakdown of the attractive forces between and and 3. Establishment of attractive forces between and (sometimes referred to as solvation or hydration)
The Dissolving and Crystallization Processes Dissolving Crystallization The Crystallization Process Involves 1.Breakdown of the attractive forces between and 2. Establishment of attractive forces between and 3. Establishment of attractive forces between and
What happens when you dissolve an ionic compound in water?? What happens when you dissolve a polar molecule in water??
What Happens When an Ionic Compound Dissolves in Water?
What Happens When a Polar Covalent Compound Dissolves in Water? dipole-dipole attractions
What happens when you try to dissolve a nonpolar molecule in water??
What happens when you try to dissolve a nonpolar molecule in water?? Non polar solvents, such as ethanol, carbon tetrachloride, ether, and hexane, are also commonly used to dissolve nonpolar solutes, such as grease and oils.
General Solubility Rule: Like Dissolves Like Polar solutes form solutions with polar solvents. Nonpolar solutes form solutions with nonpolar solvents.
Selected Polar and Nonpolar Solvents!! POLAR SOLVENTS NONPOLAR SOLVENTS water, H2O methanol, CH3OH ethanol, C2H5OH acetone, C3H6O hexane, C6H14 heptane, C7H16 toluene, C7H8 carbon tetrachloride, CCl4 methyl ethyl ketone, CH3CH2C(O)CH3 formic acid, HCOOH acetic acid, CH3COOH chloroform, CHCl3 methylene chloride, CH2Cl2 ethyl ether, CH3CH2OCH2CH3
A solution forms when these three types of attractions are of similar strength.
Solubility The is usually a limit to the solubility of one substance in another. (Gases are always soluble in each other.) Two liquids which are mutually soluble are said be miscible. The maximum solubility of a substance under a given set of conditions* is its solubility. Solubility varies with temperature and pressure (for gases).
Enthalpy of Solution The Solution Cycle solute (aggregated) + heat solute (separated) Hsolute>0, endothermic solvent (aggregated) + heat solvent (separated) Hsolvent>0, endothermic solute (separated) + solvent (separated) solution + heat Hmix<0, exothermic
Hhydration Hsoln = Hsolute + Hsolvent + Hmix Hsoln = Hsolute + Hhydration
Hsoln = Hsolute + Hsolvent + Hmix Separated components Solution
Heat of Hydration - Hlattice Hsolution Hhydration
Dissolving NaCl in Water NaCl
Dissolving NH4NO3 in Water NH 4 NO 3
Dissolving NaOH in Water NaOH
Hsolution for Various Compounds Compound Hsolution (kj/mol) KOH -57.6 LiBr -48.8 NaF +0.92 NaCl +3.9 NH4NO3 +25.7 AgNO3 +36.9
Entropy and the Solution Process Formation of a solution does not necessarily lower the potential energy of the system. For example, neon and argon mix even though there is very little difference between the attractive forces involved between molecules. Gases mix because free energy is released through the increase in the entropy of the system. Entropy is the measure of the energy dispersal of a system. Energy has a spontaneous drive to spread out over as large a volume as possible.
Entropy in the Solution Process
ΔG = ΔH - TΔS Free energy - a measure of spontaneity
Temperature Dependance of Solubility of Solids in Water Solubility is generally reported as grams of solute that will dissolve in 100 g of water. For most solids, solubility increase with an increase in temperature. Solubility curves are used to classify solutions as saturated, unsaturated, or supersaturated.
Temperature Dependance of Solubility of Gases in Water Solubility is generally reported as moles of solute that will dissolve in 1 L of solution. Because most gases are nonpolar, solubilities are generally lower than ionic compounds or polar covalent compounds. Solubility of gases decreases as temperature increases.
Temperature Dependence of Gas Solubilty in Water
Pressure Dependance of Solubility of Gases in Water - Henry s Law The solubility of a gas is directly proportional to its partial pressure. SH = khpgas kh = Henry s constant gas kh (M/atm) O2 1.3 x 10-3 N2 6.1 x 10-4 CO2 3.4 x 10-3 NH3 5.8 x 10 1 He 3.7 x 10-4
Henry s Law
Henry s Law
Pressure Dependance of Solubility of Gases in Water
Concentration of Solutions Solutions have variable composition. Therefore, descriptions of solutions must include the components and relative quantities. The terms dilute and concentrated are sometimes used, qualitatively. In general, concentration refers to the amount of solute in a given amount of solvent.
Concentration Units Molarity Molality Mole Fraction Mass Percent
Molarity (M) and Dissociation The molarity of an ionic compound allows you to determine the molarity of dissolved ions. CaCl2 (aq) = Ca 2+ (aq) + 2 Cl - (aq) A 1 M solution of CaCl2 contains: 1 mol of Ca 2+ / liter 2 mol of Cl - / liter 3 mol total of ions/liter
Molality (m) Moles of solute/1 kg of solvent Note: The definition is in terms of amount of solvent, not solution. Does not vary with temperature, since it is based on mass and not volume.
Percent Concentration
Mole Fraction (X) Mole fraction = the fraction of the moles of one component in the total moles of all components in a solution total of all the mole fractions = 1 Mole percentage = percentage of the moles of one component in the total moles of all components in a solution
Raoult s Law The vapor pressure of a volatile solvent above a solution is equal to its mole fraction of its normal vapor pressure, P o. Psolvent in solution = (XA)(P o ) Psolvent in solution is always less than P o. For a solution of a nonvolatile substance, such as salt, the boiling point of the solution is always higher than that of the pure solution.
Raoult s Law The Equilibrium Vapor Pressure of a Pure Solvent The Equilibrium Vapor Pressure of a Solution
Dissolved Solids and Vapor Pressure The effect of a dissolved solute depends on the number of solute particles. When molecular compounds dissolve, they remain as individual particles in solution. When ionic compounds dissolve, they dissociate, creating more particles. For example a 1 M solution of NaCl in water contains twice as many dissolved particles as a 1 M solution of glucose, C6H12O6.
Solutions of Volatile Solutes The Ideal Case When both the solvent and the solute can evaporate, both will be found in the vapor phase. The total vapor pressure will be the sum of the vapor pressures of the solvent and solute. Ptotal = Psolvent + Psolute The solute decreases the solvent s vapor pressure and the solvent decreases the solute s vapor pressure. Psolvent = (Xsolvent) (P 0 solvent) Psolute = (Xsolute) (P 0 solute)
Raoult s Law: Ideal Solutions In ideal solutions, the solute-solvent interactions in the solution are equal to the solvent-solvent and solute-solute interactions broken in the creation of the solution.
Raoult s Law: Non-Ideal Solutions If the solvent-solute interactions are weaker or stronger than those in the broken interactions in the pure substances, deviations from Raoult s law occur.
Colligative Properties of Solutions (a result of nonvolatile solute particles) Vapor Pressure Lowering Boiling Point Elevation (as a result of reduced vapor pressure) Freezing Point Lowering Osmotic Pressure Changes
Boiling Point Elevation For a nonvolatile solute, the boiling point of a solution is higher than that of the pure solvent. The difference between the boiling point of the solution and that of the pure solvent is directly proportional to the molal concentration of solute particles. BPsolution - BPsolvent = Tb = (m)(kb) Kb = boiling point elevation constant (units are ºC/m) (the constant is solvent dependent)
Freezing Point Depression For a nonvolatile solute, the freezing point (or melting point) of a solution is lower than that of the pure solvent. The difference between the freezing point of the solution and that of the pure solvent is directly proportional to the molal concentration of solute particles. FPsolvent - FPsolution = Tf = (m)(kf) Kf = freezing point depression constant (units are ºC/m) (the constant is solvent dependent)
Kb and Kf for Some Common Substances Substance Kb Kf Benzene 2.53 5.12 Camphor 5.95 37.7 Chloroform 3.63 4.70 Ether 2.02 1.79 Ethyl alcohol 1.22 1.99 Water 0.51 1.86
Phase diagrams of solvent and solution. A solute changes the temperature range over which a liquid remains in the liquid state.
Osmosis Osmosis - the flow of solvent through a semipermeable membrane from a solution of low concentration to a solution of high concentration Osmotic Pressure - the amount of pressure needed to prevent osmotic flow from taking place Osmotic Pressure is directly proportional to the molarity of the solute particles. = (M)(R)(T)
More opportunities exist for water molecules to cross the membrane.
isotonic solutions Equal opportunities exist for water molecules to cross the membrane from both directions.
= (M)(R)(T) =(i)(m)(r)(t) Solute nonelectrolyte i expected i measured 1 1 NaCl 2 1.9 MgSO4 2 1.3 Jacobus Henricus van 't Hoff MgCl2 3 2.7 K2SO4 3 2.6 FeCl3 4 3.4