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Chapters 12 & 13 Study Guide

by: Lauren Savage

Chapters 12 & 13 Study Guide 1442

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These notes are taken directly from the book and cover chapters 12 & 13 for the first exam. They include examples and definitions, as well as general concepts from each section.
General Chemistry 2
Dr. Rodgers
Study Guide
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This 12 page Study Guide was uploaded by Lauren Savage on Thursday September 8, 2016. The Study Guide belongs to 1442 at University of Texas at Arlington taught by Dr. Rodgers in Fall 2016. Since its upload, it has received 6 views. For similar materials see General Chemistry 2 in CHEM at University of Texas at Arlington.


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Date Created: 09/08/16
Chapter 12 Intermolecular Forces: Liquids, Solids, and Phase Changes 12.1 AN OVERVIEW OF PHYSICAL STATES AND PHASE CHANGES Phase: physically distinct, homogeneous part of a system. Intermolecular Forces: The attractive and repulsive forces among the particles—molecules, atoms, or ions—in a sample of matter.  The potential energy in the form of intermolecular forces draws molecules together.  The kinetic energy due to their random motion disperses them. It is related to their average speed and is proportional to the absolute temperature. Phase Changes: A physical change from one phase to another; usually referring to a change in physical state.  Intramolecular forces exist within each molecule. The chemical behavior of the three states is identical because all of them exist the same and are held together by identical forces.  Intermolecular forces exist between the molecules. The physical behavior of the states is different because the strengths of the forces differ depending on the state. In a gas, the PE is small compared to the KE. So, on average, the particles are far apart. Flows well. In a liquid, attractions are stronger because particles are actually touching, but they have enough KE to move and touch one another. The liquid can conform to the shape of its container. Flows slowly. In a solid, the attractions dominate so that the particles are fixed in position. Does not flow. The ability to flow influences the degree of mixing and composition in nature.  As temperature increases, the average KE does too, so the particles move faster and overcome attractions more easily.  As temperature decreases, the average KE does too, so the particles move more slowly and attractions can pull them together more easily. Gasliquid/liquidgas: as temperature drops, molecules come together for condensation. As they heat, they leave the liquid forming a vapor. Liquidsolid: as the temperature drops, the liquid freezes; when the temperature rises, fusion (melting) happens. Gassolid: deposition. Solidgas: sublimation  Exothermic: condensing, freezing, and depositing  Endothermic: melting, vaporizing, and subliming Heat (or enthalpy) of sublimation: the enthalpy change when 1 mol of a substance sublimes, and the negative of this value is the change when 1 mol of the substance deposits. Solidliquid: fusion Liquidgas: vaporization Solidgas: sublimation 12.2 QUANTATIVE ASPECTS OF PHASE CHANGES Heating-cooling curve: a plot of temperature vs time for a substance when heat is absorbed or released by the system at a constant rate. Hess’s Law: the total heat released is the sum of the heats released for the individual stages. During a phase change, heat flow occurs at a constant temperature. The Equilibrium Nature of Phase changes: 1. Open system: nonequilibrium process: some molecules at the surface have a high enough KE to overcome attractions and vaporize. Nearby molecules fill the gap, and the process continues as heat is supplied by the constant-temperature surroundings, the process continues until the entire liquid phase is gone. 2. Closed system: equilibrium process: in a vacuum, some molecules have high enough KE to vaporize, molecules in the vapor will collide with the surface, and the slower ones are attracted strongly enough to condense. Rate of vaporization: the number of molecules leaving the surface per unit time Rate of condensation: the number of molecules colliding with and entering the surface. The pressure increases and over time, the number of molecules colliding with the surface increases as the vapor becomes more populated, so the increase in pressure slows. Eventually the rate of condensation=the rate of vaporization. Then the pressure is constant. Dynamic equilibrium from liquidgasliquid Equilibrium vapor pressure: the pressure exerted by a vapor at equilibrium with its liquid in a closed system. 3. Disturbing a system at equilibrium: a. Decrease in pressure: the rate of condensation temporarily falls below the rate of vaporization because fewer molecules enter the liquid than leave it. The pressure will rise until the condensation rate increases enough for equilibrium to be reached again. b. Increase in pressure: rate of condensation temporarily exceeds the rate of vaporization because more molecules enter the liquid than leave it. The condensation rate decreases until the pressure again reaches the equilibrium value. When a system at equilibrium is disturbed, it counteracts the disturbance until it re-establishes equilibrium. The Effect of Temperature on Intermolecular Forces on Vapor Pressure The vapor pressure is affected by two factors—a change in temperature and a change in the gas itself (the type and or strength of intermolecular forces). 1. Effect of Temperature: temperature changes the fraction of molecules moving fast enough to escape the liquid or moving slowly enough to be recaptured. a. THE HIGHER THE TEMPERATURE, THE HIGHER THE VAPOR PRESSURE. 2. Effect of intermolecular forces: At a given temperature, all substances have the same KE, so molecules with a weaker intermolecular force are held together less tightly at the surface and vaporize more easily. a. THE WEAKER THE INTERMOLECULAR FORCES ARE, THE HIGHER THE VAPOR PRESSURE Vapor Pressure and Boiling Point 1. How a liquid boils: in an open container: as the temperature rises, molecules move more quickly throughout the liquid. The average KE becomes great enough in the liquid so that they form bubbles of vapor In the interior—the liquid boils. a. The boiling point is the temperature at which the vapor pressure equals the external pressure b. Once boiling begins, temperature remains the same until all the liquid is gone. 2. Effect of pressure on boiling point: boiling point is directly proportional to the applied pressure. The normal boiling point is observed at STP. Solid-Liquid Equilibria As temperature rises in a crystal, the particles move more rapidly, until some have enough KE to break free of their positions. At this point, melting begins. As more molecules become liquid, some collide with the solid and become fixed in this position again—the phases remain in contact and a dynamic equilibrium is established (when the melting rate equals the freezing rate) “MELTING POINT” the temperature remains fixed at the melting point until all the solid melts. Pressure has little effect on the rates of melting and freezing. Solid-Gas Equilibria A substance will sublime rather than melt because the intermolecular attractions are not great enough to keep the molecules near each other when they leave the solid state. Some solids do have high enough vapor pressure to sublime at ordinary conditions (Dry ice, iodine, ALL NONPOLAR MOLECULES WITH WEAK INTERMOLECULAR FORCES). Reading phase diagrams: 1. Regions of the diagrams: if a phase is placed in one of the three conditions, it is stable. 2. Lines between regions: Any point along a line shows the pressure and temperature at which the phases are in equilibrium. 3. The triple point: all three phases meet in equilibrium. 4. The critical point: the two densities become equal and the phase boundary disappears. (critical temperature and pressure). The average KE is so high that the vapor cannot be condensed at any pressure. Beyond the critical temperature—a supercritical fluid exists rather than separate liquid and gaseous phases. It expands and contracts like a goas, but has unusual solvent properties. Solid-liquid line for water: the solid form is less dense than the liquid, so water expands upon freezing. The solid-liquid line for H2O has a negative slope. 12.3 TYPES OF INTERMOLECULAR FORCES  Bonding forces are relatively strong because larger charges are closer together  Intermolecular forces are relatively weak because smaller charges are father apart How close can molecules approach each other?  Bond length—the shorter distance between two nuclei in the same molecule. One- half this distance is the covalent radius.  Van der Walls distance and radius. The longer distance is between two nonbonded atoms in adjacent molecules (called the Van der Walls distance). At this distance, intermolecular attractions balance electron-cloud repulsions—the distance is as close as one Cl2 molecule can approach another.  Van der Walls radius is one-half the closest distance between nuclei of identical nonbonded atoms. The VDW radius of an atom is always larger than its covalent radius. o VDW radii decrease across a period and increase down a group. Ion-Dipole Forces When an ion and a nearby polar molecule (dipole) attract one another. (like when an ionic compound dissolves in water). Dipole-Dipole Forces: the intermolecular attractions between oppositely charged poles of nearby polar molecules. These forces depend on the magnitude of the molecular dipole moment. For compounds of similar molar mass, the greater the molecular dipole moment, the greater the dipole-dipole forces, so the more energy it takes to separate the molecules (the BOILING POINT IS HIGHER). Smaller dipole moments boil at lower temperatures. The Hydrogen Bond: a special type of dipole force BETWEEN molecules that have an H atom bonded to a small, highly electronegative atom with lone electron pairs, specifically N, O, or F. H-N, H-O, and H-F bonds are very polar. Why N, O, and F?  The atoms are so electronegative that their covalently bonded H is highly positive  The lone pair on the N,O,or F of the other molecule can come close to the H. Why is Hydrogen Bonding significant? It is responsible for the action of many enzymes that speed metabolic reactions Polarizability and Induced Dipole Forces A nearby electric field can induce a distortion in the cloud, pulling electron density toward a positive pole of a field or pushing it away from a negative one.  For a nonpolar molecule, the distortion induces a temporary dipole moment  For a polar molecule, the distortion induces an increase in the already existing dipole moment. The source of the electric field can come from a battery, the charge of an ion, or the partial charges of a polar molecule. Polarizability: how easily the electron cloud of an atom or ion can be distorted.  Smaller particles are less polarizable than larger ones because their electrons are closer to the nucleus and therefore held more tightly. o Polarizability increases down a group because atomic size increases and larger electron clouds are easier to distory. o Polarizability decreases across a period because increasing Zeff makes the atoms smaller and holds the electrons more tightly. o Cations are less polarizable than their parent atoms because they are smaller; anions are more polarizable because they are larger.  Ion-induced dipole and dipole-induced dipole forces are the two types of charge induced dipole forces; they are most important in a solution.  Polarizability affects all intermolecular forces. London Dispersion Forces The intermolecular force responsible for the condensed states of nonpolar substances.  At any given time, there can be an instantaneous dipole (depending on electron location) o The instantaneous dipole in one atom induces a dipole in its neighbor, and they attract one another. o Dispersion forces are instantaneous dipole-induced dipole forces  They are the ONLY force between nonpolar particles, but dispersion forces contribute to the energy of attraction in all substances because they exist between ALL particles. o Except for the forces between small, highly polar molecules or molecules forming H bonds, the dispersion force is the dominant intermolecular force. The relative strength of a dispersion force depends on the polarizability of the particles, so they are weak for small particles, but stronger for larger particles. o POLARIZABILITY DEPENDS ON THE NUMBER OF ELECTRONS, WHICH CORRELATES CLOSELY WITH MOLAR MASS because heavier particles are either larger atoms or molecules with more atoms and therefore more electrons. AS MOLAR MASS INCREASES, DISPERSION FORCES INCREASE AS WELL AS BOILING POINTS. A molecule with more area allows stronger attractions (the electron cloud can be distorted). Properties of the Liquid State Surface Tension  An interior molecule is attracted by others on all sides A surface molecule is only attracted by others below and to the sides, so it experiences a NET ATTRACTION DOWNWARD To increase attractions and become more stable, a surface molecule tends to move into the interior. A liquid surface has the fewest molecules and the smallest area possible. It acts as a skin, covering the interior. Surface tension: the energy required to increase the surface area of a liquid by a given amount. The stronger the forces between particles, the more energy it takes to increase the surface area, so the greater the surface tension. Water has a high surface tension because its molecules form multiple H bonds. Surfactants decrease the surface tension of water by congregating at the surface and disrupting the H bonds. Capillarity A property that results in a liquid rising through a narrow space against the pull of gravity.  Results from a competition between the intermolecular forces within a liquid (cohesive forces) and those between the liquid and the tube wells (adhesive). Water forms a meniscus due to adhesion Mercury—cohesive forces among the mercury atoms are metallic bonds, so they are much stronger than the mostly dispersion adhesive forces between the mercury and the glass. As a result, the liquid pulls away from the walls. The surface atoms are being pulled toward the interior by mercury’s high surface tension, so the levels drop. This produces a convex meniscus. Viscosity the resistance of a fluid to flow, resulting from intermolecular attraction that impede the movement of molecules around and past one another.  Effect of temperature: VISCOSITY DECREASES WITH HEATING. Faster moving molecules overcome intermolecular forces more easily, so the resistance to flow decreases. o Think: oil becomes more fluid as it warms in a pan  Effect of molecular shape: small spherical molecules make little contact and pour easily. Long molecules make more contact and become entangled and pour slowly. Liquids consisting of longer molecules have higher viscosities. The Uniqueness of Water  The solvent power of water results from its polarity and H bonding ability Solvent Properties of Water  It dissolves ionic compounds through ion-dipole forces that separate the ions from the solid and keeps them in solution.  It dissolves polar nonionic substances, such as ethanol and glucose by H bonding.  It dissolves nonpolar atmospheric gases to a limited extent through dipole-induced dipole and dispersion forces. Thermal Properties of Water  Specific heat capacity: is higher than most other liquids. THIS MODERATES THE EARTH  Heat of vaporization: Multiple strong H bonds give water a very high heat of vaporization. Heat is converted to potential energy inside the body and breaks H bonds and evaporates sweat, resulting in a stable body temperature and minimal loss of body fluid. The sun’s energy vaporizes ocean water in warm latitudes, and the PE is released as heat to warm cooler regions when the vapor condenses to rain. This powers weather patterns. Surface properties of water  Hydrogen bonding is responsible for water’s high surface tension and high capillarity. Water has the highest surface tension of any liquid—it keeps debris resting on the surface of water, provides shelter and nutrients for fish and insects. High capillarity means water rises through the tiny spaces between soil particles, so plant roots can absorb deep groundwater during dry periods. The unusual density of solid water The hexagonal state of water leads to the open structure of ice and the snowflake shape—the large spaces within ice make the solid less dense than the liquid.  Surface ice of lakes: the ice floats and the water beneath is fluid  Nutrient turnover: during seasonal changes, the most dense layer of water on top sinks and pushes the old water on the bottom upward – this distributes nutrients and dissolved oxygen  Soil formation—freeze thaw stress cracks the rocks over eons which produces sand and soil. Structural features of solids:  Crystalline solids: well-defined shapes because of their particles—atoms, molecules, and ions occur in an orderly arrangement.  Amorphous solids: poorly defined shapes because their particles lack an orderly arrangement throughout the sample. The Crystal Lattice and the Unit Cell: the particles in a crystal are packed tightly. The crystal lattice is the regular pattern formed around a point in the center. The lattice consists of all points with identical surroundings. Unit cell: the smallest portion that gives the crystal if it is repeated in all directions. Coordination number: the number of nearest neighbors of a particle 1. In the simple cubic unit cell: the centers of 8 identical particles define the corners of a cube. 2. Body centered cubic unit cell: identical particles lie at each corner and in the center of the cube 3. Face centered cubic unit cell: identical particles lie at each corner and in the center of each face but not in the center of the cube. Particles at the corners touch those in the faces, but not each other. Chapter 13 properties of mixtures: solutions and colloids Colloid: heterogeneous mixture—two or more phases. Types of solutions: intermolecular forces and solubility: Miscible: soluble in each other in any proportion. The physical state of the solvent usually determines the physical state of the solution. Solubility: max amount that dissolves in a fixed quantity of a given solvent at a given temperature when an excess of the solute is present. Different solutes have different solubilities.  Substances with similar types of intermolecular forces dissolve in each other Types of intermolecular forces in solutions: Ion-dipole forces: principal force involved when an ionic compound dissolves in water  Forces compete causing the crystal structure to break down  Hydration shells form: as an ion separates, water molecules cluster around it. The number of water molecules in the innermost shell depends on the ion’s size.  Hydrogen bonding is the principal force in solutions of polar, O-, and N containing organic and biological compounds.  Dipole-dipole forces—in the absence of H bonding—allow polar molecules to dissolve in polar solvents.  Ion-induced dipole forces: based on polarizability: when an ion’s charge distorts the electron cloud of a nearby nonpolar molecule.  Dipole-induced dipole forces: based on polarizability: when a molecule distorts the electron cloud of a nonpolar molecule. THESE ARE WEAKER THAN ION-INDUCED DIPOLE FORCES BECAUSE THE CHARGE OF EACH POLE IS LESS THAN AN ION’S.  Dispersion forces contribute to the solubility of all solutes in all solvents, but they are the principal intermolecular force in solutions of nonpolar substances. Liquid solutions and the role of molecular polarity When the forces within the solute are similar to those within the solvent, the forces can replace each other and a solution will form.  Salts are soluble in water because of ion-dipole attractions between ion and water are similar in strength to the strong attractions between the ions and the strong H bonds between water molecules, so they can replace one another.  Salts are insoluble in hexane  Oil is insoluble in water because the weak dipole-induced dipole forces between oil and water molecules cannot replace the strong H bonds between water molecules or the extensive dispersion forces within the oil.  Oil is soluble in hexane. 1. Solubility in water is high for smaller alcohols 2. Solubility in water is low for larger alcohols Solubility in hexane is low for the smallest alcohol Solubility in hexane is high for larger alcohols.  All gases are miscible with each other.  When a gas dissolves in a solid, it occupies the spaces between the closely packed particles.  Solid-solid solutions: solids diffuse so little that their mixtures are usually heterogeneous. Some solutions can be formed by melting the solids and then mixing them and allowing them to freeze.  Alloy: mixtures of elements that have a metallic character are solid-solid solutions. INTERMOLECULAR FORCES AND BIOLOGICAL MACROMOLECULES  Polar and ionic groups attract water, but nonpolar groups do not  Distant groups on the same molecule attract one another THE STRUCTURE OF PROTEINS Proteins: a natural, linear polymer composed of any of about 20 types of amino acid monomers linked together by peptide bonds Amino acids: an organic compound with at least one carboxyl and one amino group on the same molecule; the monomer unit of a protein. Polarity of amino-acid side chain: intermolecular forces and protein shape: distant groups on the protein chain end up near each other as the chain bends.  Covalent peptide bonds create the backbone (polypeptide chain)  Helical and sheetlike segments arise from H bonds between the C=O of one peptide bonds and the N-H of another.  Polar and ionic side chains interact through dispersion forces within the nonaqueous protein interior  The –SH ends of the two cysteine side chains form a covalent –S-S- bond, a disulfide bridge, and fixes a bend.  Oppositely charged ends of ionic side chains –COO- and –NH3 groups form an electrostatic salt link that creates bend in the protein chain Soluble proteins have polar-ionic exteriors and nonpolar interiors . The amino acid sequence of a protein determines its shape, which determines its function. Soap The salt formed when a strong base reacts with a fatty acid—a nonpolar tail with a cation base. This cation base affects the properties of the soap. Nonpolar tails interact with the nonpolar grease molecules through dispersion forces Polar-ionic heads attract water molecules through ion-dipole forces and H bonds. Phospholipids have dual polarity and form a bilayer with the tails touching.  Ion dipole forces occur between polar heads and water inside and outside  Dispersion forces occur between nonpolar tails within the bilayer interior  Minimal contact exists between nonpolar tails and water Soluble proteins have polar exteriors and nonpolar interiors. They form ion-dipole and h bonding forces between water and polar groups on the exterior and dispersion forces between nonpolar groups in the interior Membrane proteins have exteriors that are partially polar and partially nonpolar. They have polar parts on the outside that goes into aqueous surroundings and nonpolar groupson the exterior portion embedded in the membrane. The Structure of DNA Nucleic acids: an unbranched polymer consisting of mononucleotides that occurs as two types: DNA and RNA, which differ chemically in the nature of the sugar portion of the mononucleotides Sugar linked to phosphate linked to sugar. Intermolecular Forces and the Double Helix  On the more polar exterior, negatively charged sugar-phosphate groups interact with the aqueous surroundings via ion-dipole forces and H bonds  In the less polar interior, flat N-containing bases stack above each other and interact by dispersion forces  Bases form specific interchain h bonds: EACH BASE IN ONE CHAIN IS ALWAYS H BONDED WITH ITS COMPLEMENTARY BASE IN THE OTHER CHAIN. The base sequence of one chain is the H bonded complement of the base sequence of the other. The total energy of the H bonds keeps the chains together, but each H bond is weak enough that a few at a time can break as the chains separate during crucial celluar processes. The Heat of Solution and its Components: 1. Solute particles separate from each other. a. Overcoming intermolecular (or ionic) attractions b. ENDOTHERMIC 2. Solvent particles separate from each other a. ENDOTHERMIC 3. Solute and solvent particles mix and form a solution. a. The different particles attract each other and come together b. EXOTHERMIC This is called the thermochemical solution cycle. Enthalpy of a solution=delta H solute+ delta H solvent+ delta H mix  If the sum of the endothermic terms (delta H solute + delta H solvent) is smaller than the exothermic term (delta H mix), the process is exothermic and delta H solution is negative.  If the sum of the endothermic terms is larger than the exothermic term, the process is endothermic and delta H solution is positive. If delta H solution is highly positive, the solute may not dissolve significantly in that solvent. Solvation: the process of surrounding a solute particle with solvent particles. Solvation in water is called hydration, so enthalpy changes for separating the water molecules and mixing the separated solute with them are combined into the heat (enthalpy) of hydration. Heat/enthalpy of hydration: Hydration of an ion is always exothermic. Heats of hydration exhibit trends based on the ion’s charge density: the ratio of the charge of an ion to its volume. The higher the charge density, the more negative delta H hydration is. The higher the charge of an ion and the smaller its radius, the closer it gets to the oppositely charged pole of an H2O molecule and the stronger the attraction. Down a group—the charge stays the same and the size increases; the charge densities decrease, as do the delta H hydration values. Across a period—ion has a smaller radius and a higher charge, so its charge density and delta H hydration are greater. For ionic compounds in water, the heat of solution is the lattice energy (always positive) plus the combined heats of hydration of the ions (always negative).  One other factor that determines whether a solute dissolves o Natural tendency of a system of particles to spread out, which results in the KE becoming more dispersed or more widely distributed.  Entropy: a thermodynamic quantity related to the number of ways the energy of a system can be dispersed through the motions of its particles. Entropy and the Three Physical States:  The more freedom of motion the particles have, the more ways they can distribute their kinetic energy; thus a liquid has higher entropy than a solid, and a goas has higher entropy than a liquid.  Thus, there is a change in entropy (delta S) associated with a phase change, and it can be positive or negative. A solution usually has a higher entropy than the pure solute and pure solvent because the number of ways to distribute the energy is related to the number of interactions between different molecules. During solution formation, systems change toward a state of lower enthalpy and higher entropy, so the sizes of delta H solution and delta S solution determine whether a solution forms.  A saturated solution is at equilibrium and contains the maximum amount of dissolved solute at a given temperature in the presence of undissolved solute. If you filter off the solution and add more solute, it doesn’t dissolve  An unsaturated solution contains less than the equilibrium concentration of dissolved solute— add more solute and more dissolves until the solution is saturated  A supersaturated solution contains more than the equilibrium concentration and is unstable relative to the saturated solution. -Most solids are more soluble at higher temperatures. -The gas solubility in water decreases with rising temperature because gases have weak intermolecular forces with water. When the temperature rises, the average KE increases, allowing the gas particles to easily overcome these forces and re-enter the gas phase.This leads to thermal pollution. -Pressure has little effect on the solubility of liquids and solids because they are almost incompressible. -the solubility of a gas is directly proportional to the partial pressure of the gas above the solution. Using henry’s law constant. S gask +H gas S=mol/L, P=atm, and Kh= mol/(L*atm) Concentration Concentration: proportion of a substance in a mixture; an intensive property—one that does not depend on the quantity of mixture. Molarity (M): amount (mol) of solute/(Volume(L) of solution)  Two drawbacks that affect its use in precise work: o Effect of temperature: a liquid expands when heated, so a unit volume of hot solution contains less solute than one of cold solution—the molarity is different o Effect of mixing: due to solute-solvent interactions that are difficult to predict, volumes may not be additive Molality (m): amount (mol) of solute/mass (kg) of solvent  Effect of temperature—mass does not change with temperature, so this is NOT a problem  Effect of mixing—masses are additive, so this is NOT a problem Parts by mass: mass of solution/mass of solution; mass percent. Parts by volume: volume of solute/volume of solution Mole fraction (X) : amount (mol) of solute/ (amount (mol) of solute + amount (mol) of solvent) To convert a term based on amount to one based on mass, need the molar mass. To convert a term based on mass to one based on volume, you need the solution density Molality includes quantity of solvent; the other terms include quantity of solution. Colligative Properties of Solutions The number of solute particles, not their chemical identity, that make the difference in physical properties.  Vapor pressure lowering  Boiling point elevation  Freezing point depression  Osmotic pressure The magnitude of a colligative property from the solute formula,which shows the number of particles in a solution and is closely related to our classification of solutes by their ability to conduct an electric current.  For strong electrolytes, the number of particles equals the number of ions in a formula unit.  Weak electrolytes form few ions  For nonelectrolytes, the number of particles equals the number of molecules. 1. Electrolytes: an aqueous solution of an electrolyte (a substance that conducts a current when it dissolves in water. A mixture of ions, in which the electrodes of an electrochemical cell are immersed, that conducts a current) conduces a current because the solute separates into ions as it dissolves. a. Strong electrolyte: soluble salts, strong acids, and strong bases— dissociate completely, so their solutions conduct well b. Weak electrolytes: weak acids and weak bases—dissociate very little, so their solutions conduct poorly 2. Nonelectrolytes: compounds such as sugar and alcohol do not dissociate into ions at all. They are nonelectrolytes because their solutions do not conduct an electric current. a. For nonelectrolytes: 1 mol of compound yields 1 mol of particles when it dissolves in solution b. For strong electrolytes: 1 mol of compound dissolves to yield the amount of ions shown in the formula unit. c. For weak electrolytes: the calculation is complicated because the solution reaches equilibrium. Nonvolatile Nonelectrolyte Solutions Contain solutes that are not ionic and thus do not dissociate, have negligible vapor pressure at the boiling point of the solvent. Vapor pressure lowering: the lowering of the vapor pressure of a solvent caused by the presence of dissolved solute particles. The vapor pressure of a nonvolatile nonelectrolyte solution is always lower than the vapor pressure of the pure solvent. 1. Why the vapor pressure of a solution is lower: the relative changes in entropies of vaporization of a solvent versus solution. A liquid vaporizes because a gas has higher entropy. In a closed container, vaporization continues until the numbers of particles leaving and entering the liquid phase per unit time are equal. Since the entropy of a solution is already higher than that of a pure solvent, so fewer solvent particles need to vaporize to reach the same entropy. With fewer particles in the gas phase, the vapor above a solution has lower pressure. 2. Quantifying vapor pressure lowering: Raoult’s Law: the vapor pressure of a solvent above a solution equals the mole fraction of solvent (X solvent the vapor pressure of the pure solvent (P solvent) Psolvent =X solvent*Psolvent An ideal solution: a solution whose vapor pressure equals the mole fraction of the solvent * vapor pressure of the pure solvent; approximated only by very dilute solutions. 3. Boiling point elevation: the increase in the boiling point of a solvent caused by the presence of dissolved solute. a solution boils at a higher temperature than the pure solvent due to vapor pressure lowering. The boiling point elevaton results because a higher temperature is needed to raise the solution’s vapor pressure to equal. 4. Freezing Point Depression: a solution freezes at a lower temperature than the pure solvent. a. Why a solution freezes at a lower T: only solvent vaporizes from solution, so solute molecules are left behind. The vapor pressure of the solution is always lower than that of the solvent, so the solution freezes at a lower temperature—only at a lower temperature will solvent particles leave and enter the solid at the same rate. b. Quantifying freezing point depression: the freezing point depression is proportional to the molal concentration of solute. 5. Osmotic pressure: this is only observed when solutions of higher and lower concentrations are separated by a semipermeable membrane—one that allows solvent, but not solute, to pass through. Osmosis: the process by which solvent flows through a semipermeable membrane from a dilute to a concentrated solution. a. Why osmotic pressure arises : when higher molecules are unable to get through themembrane, but water flows back and dilutes the other solution. Some water gets pushed back out because of the pressure difference, and this is the point of osmotic pressure—the same pressure that must be applied to prevent net movement of water from solvent to solution (or from higher to lower concentration) b. Quantifying osmotic pressure : the osmotic pressure is proportional to the number of solute particles in a given solution volume (to the molarity). Underlying themes of Colligative properties:  Each property arises because the solute particles cannot move between two phases o They cannot enter the gas phase, which leads to vapor pressure lowering and boiling point elevation. o They cannot enter the solid phase, which leads to freezing point depression o They cannot cross a semipermeable membrane, which leads to osmotic pressure.  The presence of solute decreases the mole fraction of solvent, which lowers the number of solvent particles leaving the solution per unit time. This lowering maintains higher entropy and requires a new balance in numbers of particles moving between phases per unit time. Particle size plays a defining role in three types of mixtures:  Suspensions: a heterogeneous mixture containing particles that are distinct from the surrounding medium.  Solutions: a homogenous mixture in which the particles are invisible  Colloids: colloidal dispersions: a heterogeneous mixture in which a solute-like phase is dispersed throughout a solvent-like phase. The particles are larger than simple molecules but too small to settle out. 1. Particle size and surface area: a colloid has a very large total surface area. The large surface area of the colloid attracts other particles through various intermolecular forces. 2. Classification of colloids: 3. Tyndall effect and Brownian motion: the scattering of light by a colloid. The dispersed particles have sizes similar to the wavelengths of visible light. (dust in sunlight) 4. Stabilizing and destabilizing colloids: colloidal particles dispersed in water have charged surfaces that stabilize the colloid through ion-dipole forces. Molecules with dual polarities form spherical micelles . Water softening via Ion exchange  Hard water: water with large amounts of 2+ ions such as Ca2+ and Mg+  Water softening: the process of replacing the hard-water ions Ca2+ and Mg2+  Ion exchange system: a process of softening water by exchanging one type of ion for another by binding the ions on a specifically designed resin. Membrane Processes and Reverse Osmosis Reverse osmosis: a process for preparing drinkable water that uses an applied pressure greater than the osmotic pressure to remove ions from an aqueous solution, typically seawater.


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