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Properties of Solutions Brown, LeMay Ch 13 AP Chemistry Edited significantly LJC 2015 CaCl2 (aq) 1 13.1: Types of Solutions Example Air (g in g) Soda (g in l) H2 in Pt (g in s) Alcoholic beverages (l in l) Solvent Solute Sea water (s in l) Brass (s in s) 3 13.1: Types of Solutions Example Air (g in g) Soda (g in l) H2 in Pt (g in s) Alcoholic beverages (l in l) Solvent N2 Solute O2 H2O Pt CO2 H2 H2O C2H5OH Sea water (s in l) H2O Brass (s in s) Copper (55% – 90%) NaCl (one of many salts) Zinc (10% – 45%) 4 Solutions • Solutions are homogeneous mixtures of two or more pure substances. • In a solution, the solute is dispersed uniformly throughout the solvent. Solutions How does a solid dissolve into a liquid? What ‘drives’ the dissolution process? What are the energetics of dissolution? How Does a Solution Form? 1. Solvent molecules attracted to surface ions. 2. Each ion is surrounded by solvent molecules. 3. Enthalpy (DH) changes with each interaction broken or formed. Ionic solid dissolving in water How Does a Solution Form? 1. Solvent molecules attracted to surface ions. 2. Each ion is surrounded by solvent molecules. 3. Enthalpy (DH) changes with each interaction broken or formed. How Does a Solution Form? The ions are solvated (surrounded by solvent). If the solvent is water, the ions are hydrated. What intermolecular force is present? How Does a Solution Form? The ions are solvated (surrounded by solvent). If the solvent is water, the ions are hydrated. What intermolecular force is present? ION-DIPOLE Energy Changes in Solution To determine the enthalpy change, we divide the process into 3 steps. 1. 2. 3. Energy Changes in Solution To determine the enthalpy change, we divide the process into 3 steps. 1. Separation of solute particles. 2. Separation of solvent particles to make ‘holes’. 3. Formation of new interactions between solute and solvent. Energy Changes in Solution Predict: which of these steps would likely be endothermic? Which would be exothermic? Enthalpy Changes in Solution Start The enthalpy change of the overall process depends on DH for each of these steps. Enthalpy changes during dissolution DHsoln = DH1 + DH2 + DH3 The enthalpy of solution, DHsoln, can be either positive or negative. DHsoln (MgSO4)= -91.2 kJ/mol --> exothermic DHsoln (NH4NO3)= 26.4 kJ/mol --> endothermic Why do endothermic processes sometimes occur spontaneously? Some processes, like the dissolution of NH4NO3 in water, are spontaneous at room temperature even though heat is absorbed, not released. Enthalpy Is Only Part of the Picture There is also ENTROPY. Entropy is a measure of: • Dispersal of energy in the system. • Number of microstates (arrangements) in the system. • OR, put simply, degree of disorder (more on this in chap 19) Which beaker is the favored energy state? Enthalpy Is Only Part of the Picture There is also ENTROPY. Entropy is a measure of: • Dispersal of energy in the system. • Number of microstates (arrangements) in the system. • OR, put simply, degree of disorder (more on this in chap 19) (b) has greater entropy, is the favored state Entropy changes during dissolution What were the three steps of solution formation? 1. 2. 3. What would the entropy change during each step be? SAMPLE EXERCISE 13.1 Assessing Entropy Change In the process illustrated below, water vapor reacts with excess solid sodium sulfate to form the hydrated form of the salt. The chemical reaction is Does the entropy of the system increase or decrease? Dissolution vs reaction Ni(s) + HCl(aq) NiCl2(aq) + H2(g) dry NiCl2(s) • Dissolution is a physical change—you can get back the original solute by evaporating the solvent. • If you can’t, the substance didn’t dissolve, it reacted. Degree of saturation • Saturated solution Solvent holds as much solute as is possible at that temperature. Undissolved solid remains in flask. Dissolved solute is in dynamic equilibrium with solid solute particles. Degree of saturation • Unsaturated Solution Less than the maximum amount of solute for that temperature is dissolved in the solvent. No solid remains in flask. Degree of saturation • Supersaturated Solvent holds more solute than is normally possible at that temperature. These solutions are unstable; crystallization can often be stimulated by adding a “seed crystal” or scratching the side of the flask. Degree of saturation How can you tell whether a solution is unsaturated, saturated or supersaturated? How much solute can be dissolved in a solution? SOLUBILITY is how much solute can be dissolved in a solution. More on this in Chap 17 (solubility products, p 739) Factors Affecting Solubility What do you observe? How does this illustrate “Like Dissolves Like”? Factors Affecting Solubility Example: ethanol in water Ethanol = CH3CH2OH The stronger the intermolecular attractions between solute and solvent, the more likely the solute will dissolve. Intermolecular forces = H-bonds; dipole-dipole; dispersion Ions in water also have ion-dipole forces. Factors Affecting Solubility Which would be more soluble in water: cyclohexane or glucose? Factors Affecting Solubility Glucose (which has hydrogen bonding) is very soluble in water. Cyclohexane (which only has dispersion forces) is not water-soluble. Factors Affecting Solubility • Which vitamin is fat soluble? • Which vitamin is water soluble? Factors Affecting Solubility • Vitamin A is soluble in nonpolar compounds (like fats). • Vitamin C is soluble in water. Which vitamin is water-soluble and which is fat-soluble? Gases in Solution What do you observe about the solubility of gases in water? Why? (Be CL.EV.ER) Gases in Solution • In general, the solubility of gases in water increases with increasing mass. Why? • Larger molecules have stronger dispersion forces. Gases in Solution Increasing pressure above solution forces more gas to dissolve. • The solubility of liquids and solids does not change appreciably with pressure. • But, the solubility of a gas in a liquid is directly proportional to its pressure. Henry’s Law Sg = kPg where • Sg is the solubility of the gas; • k is the Henry’s law constant for that gas in that solvent; • Pg is the partial pressure of the gas above the liquid. Temperature What do you observe about the solubility of solids as T increases? Temperature • What do you observe about the solubility of gases as T increases? • Is this the same or different from what you observed for solids? Carbonated soft drinks are more “bubbly” if stored in the refrigerator. Warm lakes have less O2 dissolved in them than cool lakes. Practice: what IMF’s are involved in making this solution? NaCl (s) → Na+ (aq) + Cl- (aq) Practice: what IMF’s are involved in making this solution? NaCl (s) + H2O (l) → Na+ (aq) + Cl- (aq) • Ion-dipole interactions > H-bonds (H2O···H2O) < Ionic bonds (Na+ Cl-) • The increase in disorder also drives the dissolving process. http://phet.colorado.edu/en/simulation/soluble-salts Ways to Express Concentration (Card #2) • Mass Percent • No longer on AP Exam! mass solute % 100 mass solute mass solvent • Mole Fraction: commonly used for gases mol A XA total mol • Molarity: commonly used for solutions Varies with T mol solute M Liters solution • Molality: commonly used for colligative properties Does not vary with T No longer on AP exam! mol solute m kg solvent 42 Solubility Vocabulary (Card #3) http://www.youtube.com/watch?v=VTmfQUNLlMY Saturated: a solution that is in equilibrium with undissolved solute (appears as solution and crystals) Solubility: the amount of solute needed to form a saturated solution Unsaturated: a solution containing less than the saturated amount (appears as solution only) Supersaturated: a solution containing more than the saturated amount, yet appears unsaturated. 43 Solubility Solubility: go to the temperature and up to the desired line, then across to the Yaxis. This is how many g of solute are needed to make a saturated solution of that solute in 100g of H2O at that particular temperature. At 40oC, the solubility of KNO3 in 100g of water is 64 g. In 200g of water, double that amount. In 50g of water, cut it in half. Supersaturated If 120 g of NaNO3 are added to 100g of water at 30oC: 1) The solution would be SUPERSATURATED, because there is more solute dissolved than the solubility allows 2) The extra 25g would precipitate out 3) If you heated the solution up by 24oC (to 54oC), the excess solute would dissolve. Unsaturated If 80 g of KNO3 are added to 100g of water at 60oC: 1) The solution would be UNSATURATED, because there is less solute dissolved than the solubility allows 2) 26g more can be added to make a saturated solution 3) If you cooled the solution down by 12oC (to 48oC), the solution would become saturated 13.4: Factors Affecting Solubility (Card #4) 1. “Like dissolves like.” • Miscible: liquids that mix (polar or ionic solute with polar solvent, or nonpolar with nonpolar) • Immiscible: liquids that do not mix (polar or ionic solute with nonpolar solvent) • Covalent network solids do not dissolve in polar or nonpolar solvents. 47