Properties of Solutions

Properties of Solutions Chapter 13 Chapter 13

The Solution Process • A solution is a homogeneous mixture of solute (present in smallest amount) and solvent (present in largest amount). • Solutes and solvent are components of the solution. • In the process of making solutions with condensed phases, intermolecular forces become rearranged. Chapter 13

The Solution Process • Consider NaCl (solute) dissolving in water (solvent): • the water H-bonds have to be interrupted, • NaCl dissociates into Na+ and Cl-, • ion-dipole forces form: Na+…-OH2 and Cl- …+H2O. • We say the ions are solvated by water. • If water is the solvent, we say the ions are hydrated. Chapter 13

The Solution Process • Chemical Process • There are solutions that form by chemical processes. • Consider: • Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g). • Note the chemical form of the substance being dissolved has changed (Ni  NiCl2). • When all the water is removed from the solution, no Ni is found only NiCl2.6H2O. Therefore, Ni dissolution in HCl is a chemical process. Chapter 13

The Solution Process • Physical process • Example: • NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq). • When the water is removed from the solution, NaCl is found. Therefore, NaCl dissolution is a physical process. Chapter 13

Ways of Expressing Concentration • All methods involve quantifying amount of solute per amount of solvent (or solution). • Generally amounts or measures are masses, moles or liters. • Qualitatively solutions are dilute or concentrated. • Definitions: Chapter 13

Ways of Expressing Concentration • Parts per million (ppm) can be expressed as 1 mg of solute per kilogram of solution. • If the density of the solution is 1g/mL, then 1 ppm = 1 mg solute per liter of solution. • Parts per billion (ppb) are 1 g of solute per kilogram of solution. Chapter 13

Ways of Expressing Concentration • Mole Fraction, Molarity, and Molality • Recall mass can be converted to moles using the molar mass. • Recall • Recall Chapter 13

Ways of Expressing Concentration • Mole Fraction, Molarity, and Molality • We define • Converting between molarity (M) and molality (m) requires density. Chapter 13

Saturated Solutions and Solubility • Terms • Dissolve: solute + solvent  solution. • Crystallization: solution  solute + solvent. • Saturation: crystallization and dissolution are in equilibrium. • Solubility: amount of solute required to form a saturated solution. • Supersaturated: a solution formed when more solute is dissolved than in a saturated solution. Chapter 13

The Solution Process • Energy Changes and Solution Formation • There are three energy steps in forming a solution: • separation of solute molecules (H1), • separation of solvent molecules (H2), • formation of solute-solvent interactions (H3). • Define the enthalpy change in the solution process as • Hsoln = H1 + H2 + H3. • Hsoln can either be positive or negative depending on the intermolecular forces. • NaOH added to water: Hsoln = -44.48 kJ/mol. • NH4NO3 added to water: Hsoln = + 26.4 kJ/mol Chapter 13

The Solution Process • Energy Changes and Solution Formation • Breaking attractive intermolecular forces is always endothermic. • Forming attractive intermolecular forces is always exothermic. • To determine whether Hsoln is positive or negative: • H1 and H2 are both positive. • H3 is always negative. • It is possible to have either H3 > (H1 + H2) or H3 < (H1 + H2). Chapter 13

The Solution Process Energy Changes and Solution Formation Chapter 13

The Solution Process • Energy Changes and Solution Formation • For some examples: • NaCl in gasoline: the ion-dipole forces are weak because gasoline is non-polar. Therefore, the ion-dipole forces do not compensate for the separation of ions. • Water in octane: water has strong H-bonds. There are no attractive forces between water and octane to compensate for the H-bonds. Chapter 13

The Solution Process • Solution Formation, entropy and disorder • Hsoln can predict solubility in part – if it is too endothermic a solution will not form. PbS (E-13) • But how to explain endothermic process of solution? • Another part comes from the concept of “entropy” • Entropy is related to disorder, generally speaking, “disorder” is a general trends, only subject to energy change. • Solution process is a process of increase disorder. Chapter 13

The Solution Process • Solution Formation, Spontaneity, and Disorder • Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids. Therefore, they spontaneously mix even though Hsoln is very close to zero. Chapter 13

Factors Affecting Solubility • Solute-Solvent Interactions • Polar liquids tend to dissolve in polar solvents. • Miscible liquids: mix in any proportions. • Immiscible liquids: do not mix. • Intermolecular forces are important: water and ethanol are miscible because the broken hydrogen bonds in both pure liquids are re-established in the mixture. Chapter 13

Factors Affecting Solubility • Solute-Solvent Interactions • The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water. • The number of -OH groups within a molecule increases solubility in water. Chapter 13

Factors Affecting Solubility • Solute-Solvent Interactions • Generalization: “like dissolves like”. • The more polar bonds in the molecule, the better it dissolves in a polar solvent. • The less polar the molecule the less it dissolves in a polar solvent and the better is dissolves in a non-polar solvent. • Network solids do not dissolve because the strong intermolecular forces in the solid are not re-established in any solution. Chapter 13

Factors Affecting Solubility • Pressure Effects on gas • Solubility of a gas in a liquid is a function of the pressure of the gas. • The higher the pressure, the more molecules of gas are close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution. • Therefore, the higher the pressure, the greater the solubility. • The lower the pressure, the fewer molecules of gas are close to the solvent and the lower the solubility. Chapter 13

Factors Affecting Solubility Pressure Effects Chapter 13

Factors Affecting Solubility • Pressure Effects • Henry’s Law: • Cg is the solubility of gas, Pg the partial pressure, k = Henry’s law constant. • Carbonated beverages are bottled under > 1 atm. As the bottle is opened, decreases and the solubility of CO2 decreases. Therefore, bubbles of CO2 escape from solution. Chapter 13

Factors Affecting Solubility • Temperature Effects • Experience tells us that sugar dissolves better in warm water than cold. • As temperature increases, solubility of solids generally increases. • Very few, solubility decreases as temperature increases (e.g. Ce2(SO4)3). Chapter 13

Factors Affecting Solubility Temperature Effects Chapter 13

Factors Affecting Solubility • Temperature Effects on Gas solute • Experience tells us that carbonated beverages go flat as they get warm. Chapter 13

Factors Affecting Solubility • Temperature Effects • Experience tells us that carbonated beverages go flat as they get warm. • Gases are less soluble at higher temperatures. • Thermal pollution: if lakes get too warm, CO2 and O2 become less soluble and are not available for plants or animals. Chapter 13

Colligative Properties - properties depend on quantity of solute molecules. Chapter 13

Colligative Properties • Lowering the Vapor Pressure • Non-volatile solvents reduce the ability of the surface solvent molecules to escape the liquid. • Therefore, vapor pressure is lowered. • The amount of vapor pressure lowering depends on the amount of solute. • Raoult’sLaw - • PA is the vapor pressure with solute, PA is the vapor pressure without solvent, and A is the mole fraction of A. Chapter 13

Colligative Properties • Raoult’s Law • Ideal solution: one that obeys Raoult’s law. • Raoult’s law breaks down when the solvent-solvent and solute-solute intermolecular forces are different from solute-solvent intermolecular forces. Chapter 13

Colligative Properties • Boiling-Point Elevation • Non-volatile solute lowers the vapor pressure. • At 1 atm (normal boiling point of pure liquid) there is a lower vapor pressure of the solution. Therefore, a higher temperature is required to teach a vapor pressure of 1 atm for the solution (Tb). • Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m: Chapter 13

Colligative Properties • Freezing-Point Depression • When a solution freezes, pure solvent solid is formed first, as solute molecules do not fit the solvent crystal. • The solid consists the solvent only, the nucleation is more difficult than in pure solvent. The temperature must be lower. • If solute is in low concentration, decrease in freezing point (Tf) is directly proportional to molality • Kfis the molal freezing-point-depression constant. Chapter 13

Colligative Properties Boiling-Point Elevation On the phase diagram, the triple point occurs at a lower temperature because of the lower vapor pressure and lower freezing point for the solution. Chapter 13

Colligative Properties • Osmosis • Semipermeable membrane: permits passage of some components of a solution. Example: cell membranes • Osmosis: the movement of a solvent from low solute concentration to high solute concentration, in order to equalize the solute concentrations on the two sides • Osmosis can be measured as pressure - Chapter 13

Colligative Properties • Osmosis • As solvent moves across the membrane, the fluid levels in the arms becomes uneven. • Eventually the pressure difference between the arms stops osmosis. • Osmotic pressure, , is the pressure required to stop osmosis Chapter 13

Colligative Properties • Osmosis • Isotonic solutions: two solutions with the same  separated by a semipermeable membrane. • Hypotonic solution: with a lower concentration of solute, higher osmotic pressure. • Hypertonic solution: with a higher concentration of solute, lower osmotic pressure. Chapter 13

Colligative Properties • Osmosis - Crenation and Hemolysis • Hemolysis is when a red blood cell is placed in a hypertonic solution and the cell burst because water is coming in. • Crenation is when a red blood cell is placed in a hypotonic solution and the cell shrinks because water is coming out. Chapter 13

Colloids • Colloids are suspensions in which the suspended particles are larger than molecules but too small to drop out of the suspension due to gravity. • Particle size: 10 to 2000 Å. • Types of colloid: • aerosol (gas + liquid or solid, e.g. fog and smoke), • foam (liquid + gas, e.g. whipped cream), • emulsion (liquid + liquid, e.g. milk), • sol (liquid + solid, e.g. paint), • solid foam (solid + gas, e.g. marshmallow), • solid emulsion (solid + liquid, e.g. butter), • solid sol (solid + solid, e.g. ruby glass). Chapter 13

Colloids • Tyndall effect: ability of a colloid to scatter light. The beam of light can be seen through the colloid. • Scattering happens when the cross-section of an individual particulate is the range of roughly between 40 and 900 nanometers, i.e., somewhat below or near the wavelength of visible light (400–750 nanometers). Chapter 13

Colloids • Hydrophilic and Hydrophobic Colloids • Focus on colloids in water. • “Water loving” colloids: hydrophilic. • “Water hating” colloids: hydrophobic. • Molecules arrange themselves so that hydrophobic portions are oriented towards each other. • If a large hydrophobic macromolecule (giant molecule) needs to exist in water (e.g. in a cell), hydrophobic molecules embed themselves into the macromolecule leaving the hydrophilic ends to interact with water. Chapter 13

Colloids • Hydrophilic and Hydrophobic Colloids • Typical hydrophilic groups are polar (containing C-O, O-H, N-H bonds) or charged. • Hydrophobic colloids need to be stabilized in water. • Adsorption: when something sticks to a surface we say that it is adsorbed. • If ions are adsorbed onto the surface of a colloid, the colloids appears hydrophilic and is stabilized in water. • Consider a small drop of oil in water. • Add to the water sodium stearate. Chapter 13

Colloids Hydrophilic and Hydrophobic Colloids Chapter 13

Colloids • Hydrophilic and Hydrophobic Colloids • Sodium stearate has a long hydrophobic tail (CH3(CH2)16-) and a small hydrophobic head (-CO2-Na+). • The hydrophobic tail can be absorbed into the oil drop, leaving the hydrophilic head on the surface. • The hydrophilic heads then interact with the water and the oil drop is stabilized in water. Chapter 13

Colloids Hydrophilic and Hydrophobic Colloids Chapter 13

Colloids • Hydrophilic and Hydrophobic Colloids • Most dirt stains on people and clothing are oil-based. Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by stabilizing the colloid in water. • Emulsifying agents help form an emulsion. Chapter 13

Colloids • Removal of Colloidal Particles • Colloid particles are too small to be separated by physical means (e.g. filtration). • Colloid particles are coagulated (enlarged) until they can be removed by filtration. • Methods of coagulation: • heating (colloid particles move and are attracted to each other when they collide); • adding an electrolyte (neutralize the surface charges on the colloid particles). • Dialysis: using a semipermeable membranes separate ions from colloidal particles. Chapter 13

13.14, 28, 48, 60, 62, 70, 72, 82, 100, 104, 108. Chapter 13

Properties of Solutions