Unit 3 Material
Unit 3 Material CHM 1220
Popular in General Chemistry 1
Popular in Chemistry
This 15 page Study Guide was uploaded by Necromancer23 on Monday October 19, 2015. The Study Guide belongs to CHM 1220 at Wayne State University taught by Maryfrances Barber in Summer 2015. Since its upload, it has received 53 views. For similar materials see General Chemistry 1 in Chemistry at Wayne State University.
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Date Created: 10/19/15
Chapter 11 States of Matter Liquids and Solids 111 Comparison of Gases Liquids and Solids I Gases Compressiblefluids a Composed of moleculesatoms in constant random motion through mostly empty space i Mostly empty space makes gas easily compressible relates to VolumePressure relationship b Fluid molecules move easily relative to another Liquids relatively incompressible fluids a Composed of tightly packed molecules that are in constant random motion i Lack of empty space makes it nearly incompressible but still a fluid state I Solids nearly incompressible rigid a Made up of atoms molecules or ions b Exist in close contact oscillate vibrate about fixed sites i Makes rigid c Compact structure makes it incompressible IV Ideal Gas Law PVnRT a Gases closely follow PVnRT 1 P is pressure 2 V is pressure 3 n is number of moles 4 R is a gas law constant 5 T is temperature must convert to Kelvin Celsius 27315 ii Simplicity is results from nearly negligible forces of interaction between molecules and nearly negligible molecular size compared with total volume of gas 1 Still should attempt to account for these factors by using the quotreal gas formula Pn2av2VnbnRT a Discussed further in chapter 5 b No equation for solids or liquids i Most important part is state changes V Changes of StatePhase transition a A change of a substance from one state to another b Phase homogenous part of a system given state of either a substance or a solution 112 Phase Transitions Phases 0 MeltingFusion change of a solid to a liquid b Freezing change of a liquid to the solid state c Vaporization change of a liquid to the vapor d Sublimation change of a solid to vapor i Vapor pressure of solids is quite low and sometimes appreciable ii You can reduce the pressure very low to raise the vapor pressure so you can purify solids that easily vaporize e f g 1 Examples freezedrying food ground coffee Condensation change of a gas to a liquid Deposition change of a gas to a solid Liquefaction change of a substance that is normally gas to a liquid state Vapor Pressure a b Liquids and some solids are continuously vaporizing Vapor Pressure partial pressure of the vapor over the liquid measured at equilibrium at a given temperature i Molecules in a liquid have a distribution of kinetic energies molecules moving away from the surface and toward vapor phase escape only if kinetic energies are greater than a certain minimum value equal to the potential energy from the attraction of molecules in the body of the liquid 1 Gain sufficient energy through molecular collisions Equilibrium when rates of vaporization and condensation are equal vapor pressure is steady i Measure vapor pressure here Dynamic Equilibrium equilibrium in which molecular processes are continuously occurring represented by a double arrow Volatile liquids and solids with a high vapor pressure at normal temperatures Boiling and Melting Point a b c d l Temperature Boiling Point temperature at which the vapor pressure of a liquid equals the pressure exerted on the liquid atmospheric pressure atm unless otherwise noted i As temperature increases vapor pressure increases and when vapor pressure equals the atmospheric pressure stable bubbles of vapor form within the liquid this is boiling ii Normal boiling point boiling point at 1 atm because pressure exerted can vary boiling point can vary iii As external pressure increases boiling point increases iv Higher altitude lower pressure v The lower the pressure the lower the temperature to boil Freezing Point temperature at which a pure liquid changes to a crystalline solid Melting point temperature at which a crystalline solid changes to a liquid i Occurs at same temperature as freezing ii Melting points are affected noticeably only by large pressure changes Melting and boiling points can be used to identify a substance i Characteristic physical property FreezingMelting Point Intermolecular Bonds Breaking 2 Liquid and Gas 0 Vaporization Equilibrium Vapor Pressure Measured Here Boiling Point V Heat Added V Heat Phase Transition a QsomoAT i S is the specific heat ii m is mass iii AT is the change in temperature b Every change of state requires a gain or loss of energy as heat The flat region in phase change graphs i Heat is added at a constant rate so the length of each flat region is proportional to the heat of phase transition d HeatEnthalpy of Fusion AHfus the heat needed to melt a solid i Q moles o AHfus e HeatEnthalpy of Vaporization AHvap the heat needed to vaporize a liquid i More heat needed for vaporization than melting because in vaporization you need to break intermolecular forces while melting only needs enough energy to allow molecules to escape their rigid structure 1 AHvapgtAHfus 2 Q moles o AHvap V ClausiusClapeyron Equation Relating Vapor Pressure and Liquid Temperature a Approximate equation for the variation of the logarithm of the vapor pressure or solid with absolute temperature i lnPATB 1 natural logarithm of vapor pressure is lnP 2 A and B are positive constants ii Plotting lnP versus 1T will result in a straight line with a slope of A 1 YAxB a YlnP b X1T 2 Derived from thermodynamics assuming the gas behaves in an ideal behavior iii Results in ClausiusClapeyron Equation 1 Shows A is equal to AHvap a lnP AHvapRTB 2 Two point version useful for ideal gas calculations a lnP2lnP1AHvapRT2AHvapRT1BB b InPzPl AHvapR1T1391T2 113 Phase Diagrams Phase diagrams graphical way to summarize conditions under which the different states of a substance are stable a Divided into 3 regions solid liquid and gas i In each region the state is solid unless on the curve in which case the two states are in equilibrium b Triple Point point on a phase diagram representing the temperature and pressure at which three phases of a substance coexists in equilibrium i Can have multiple ii Water s triple point defines the Kelvin temperature scale Melting Point Curve AB a Nearly vertical due to being only slightly affected by pressure i If a liquid is more dense than the solid like waterice in which ice floats the melting point decreases with pressure and leans slightly left ii If the liquid is less dense than the solid the curve leans slightly right III VaporPressure Curves for the Liquid and the Solid AC AD a AC gives vapor pressure and boiling point of liquid b AD gives vapor pressure of a solid IV Critical Temperature and Pressure a Meniscus welldefined border between liquid and solid b Supercritical fluid temperature and pressure above where meniscus disappears i Not liquid or gas phase 3 c Critical Temperature temperature above which the liquid state of a substance no longer exists regardless of pressure d Critical Pressure vapor pressure at critical temperature minimum pressure that must be applied to a gas at the critical temperature to liquefy it e Critical Point point where both temperature and pressure have their critical values H20 s Curve Supercritical Fluid Area Liquid Solid atm Gas As pressure increases you cross from gas to solid to liquid Temp in Kelvin AD sublimation AB melting AC vaporization A is Triple point 114 Properties of Liquids Surface Tension and Viscosity I Vapor Pressure is the Equilibrium partial pressure of this vapor over the liquid a Increases with temperature II Boiling Point temperature at which vapor pressure equals pressure applied to liquid III Important properties of Liquids boiling point vapor pressure surface tension viscosity a All depend on intermolecular forces IV Surface Tension a Net forcesattraction i Inside body of liquid attracted equally in all directions so no net force ii At surface experiences net attraction towards interior molecules resulting in a reduction of surface area 1 Explains spherical rain drops spheres are the smallest surface area for any geometrical shape b Surface Tension energy required to increase the surface area of a liquid by a unit amount i Can be affected by dissolved substances ii Capillary tension rise in water surface in test tube happens because water molecules are attracted to glass so it rises and the level rises to reduce surface tension 1 Final water level is a balance of surface tension and potential energy to resist gravity 2 Meniscus has edges curved upwards V Viscosity resistance to flow that is exhibited by all liquids and gases a Measurement in liquids i time for a given amount to flow through a capillary tube ii time it takes for a steel ball of a given radius to fall through a column of liquid 115 Intermolecular Forces Explaining Liquid Properties I Introduction a Intermolecular forces forces of interaction between molecules weak attraction i Explains physical properties of liquids ii Three types between natural molecules b Van der Waals forces general term for intermolecular forces that include dipoledipole and London forces II DipoleDipole Forces a Attraction intermolecular force resulting from the tendency of polar molecules to align themselves such that the position is positive end of one molecule near the negative end of another b Polar molecules have dipole moments due to the electronic structure of the molecule i Alignment creates net attraction between molecules III London Dispersion Forces a Instantaneous Dipoles at some instant there are more electrons on one side than the other i Two opposite instantaneous dipoles attract each other and electrons of both move in sync b London Forces dispersion weak attractive forces between molecules resulting from the small instantaneous dipoles that occur because of the varying positions of the electrons during their motion about the nuclei i Increase with molecular mass due to more electrons and larger atoms so they are easier to polarize 1 More easy to distort into instantaneous dipoles because electrons move further from the nuclei ii More compact means less London forces and a lower melting point IV Intermolecular Forces and the Properties of Liquids a Vapor Pressure of a liquid depend on intermolecular forces i Strong intermolecular forces means a low vapor pressure b London forces are always present and usually dominant DipoleDipole only in small polar molecules or large molecules with large dipole moments i Represent a small portion of the active intermolecular forces at play think 20 d Normal boiling points are proportional to the energy of intermolecular attraction i Low boiling point means weak intermolecular forces e Surface tension increases with strength of attractive forces and molecular mass because London forces are dominant f Viscosity partly depends on intermolecular forces i Increase in attractive forces means an increased resistance to flow ii Periodic trends increases from top to bottom of column iii High viscosity happens with tangling of molecules long molecules means a high viscosity V Hydrogen Bonding weak to moderate attractive force that exists between a hydrogen atom covalently bonded to a hydrogen atom covalently bonded to a very electronegative atom X and a lone pair of electrons on another small electronegative atom a Electronegative elements F fluoride 0 oxygen N nitrogen b X H Y 116 Classifications of Solids by Type of Attraction of Units I Introduction a Solids i Consists of structural units atoms molecules and ions that are attracted to one another strongly enough to create a rigid substance ii Classified by force holding structural units together 1 Intermolecularforces 2 Chemical bonds metallic ionic covalent iii Four types molecular metallic ionic covalent II Types of Solids a Molecular solid that consists of atoms or molecules held together by intermolecular forces b Metallic solid that consists of positive cores of atoms held together by a surrounding quotseaquot of electrons i Positively charged atomic cores are surrounded by delocalized electrons c Ionic solid that consists of cations and anions held together by the electrical attraction of opposite charges ionic bonds d Covalent Network Solid solid that consists of atoms held together in large networks or chains by covalent bonds i Example diamonds III Physical properties a For a solid to melt the forces holding the structural units in their sites must be overcome i Molecular solids weak forces means a low melting point ii IonicCovalent Network Solids chemical bonds must be broken melting point is higher 1 Melting point increases with lattice energy a Lattice energy is the energy needed to separate a crystal into isolated ions in the gaseous state represents strength of attraction of ions in the solid iii Metallic high melting point but high variability 1 IA and IIA have low melting points but increase as you move right into transition metals until IIB when the melting point becomes low again b Hardness and Structure i Hardness how easily structural units of a solid can be moved relative to one another and strength of attractive forces between units 1 Weak forces means the substance is soft ii Molecular and Ionic crystals are brittle fracture along crystal planes iii Metallic crystals are malleable can be shaped by a hammer c Electrical conductivity and Structure i Metals have good conductivity ii CovalentIonic Solids are nonconductors 1 Ionic conduct in liquid state Chapter Five The Gaseous State 50 Introduction I When variables are directly proportional their ratio is constant II When variables are inversely proportional their product is constant III Gases a Compressible into smaller volumes b Can relate pressure volume temperature and molar amounts of a substance through the ideal gas law c Composed of molecules in constant motion i Kineticmoleculartheory 51 Gas Pressure and Measurement I Pressure force exerted per unit of surface a Force mass 0 constant acceleration of gravity b Pressure force area II Acceleration change of speed per unit time III SI Units a Acceleration ms2 b Pressure pascal Pa or kgms2 IV NonSI Units a Millimeters of Mercury mmHg or torr unit of pressure equal to that exerted by a column of mercury 1mm high at 000 C b Atmosphere atm unit of pressure equal to exactly 760 mmHg V Measurement devices a Barometer device for measuring pressure of atmosphere height measures pressure b Manometer device that measures the pressures of a gas or liquid in a vessel c In a barometermanometer Pgdh gdh in above units 52 Empirical Gas Laws I Boyle s Law Relating Volume and Pressure CHANGE IN PRESSURE a Compressibility property of gas ability to be squeezed into a smaller volume by applying pressure b Boyle s Law volume of a sample of gas at a given temperature is inversely proportional to the applied pressure i PVconstant ii Used to calculate volume when pressure changes iii PressureVolume product not precisely constant deviate at high pressure II Charle s Law Relating Volume and Temperature CHANGE IN TEMPERATURE a Temperature affects volume volume increases when temperature increases and reverse i Sample of gas at a fixed pressure will increase linearly in volume with temperature b Charle s Law the volume occupied by any sample of gas at a constant pressure is directly proportional to the absolute temperature i VTconstant ll VfViXTf Ti iii Derived from intersection of lines at 27315 put into the equation Vbta or a27315b or VbT 1 Only when T in Kelvins iv Gases deviate at high pressure and low temperature III Combined Gas Law Relating Temperature Volume and Pressure a Combination of Boyle s and Charle s law the volume of a given amount of gas is proportional to the absolute temperature divided by the pressure b Can be written i Vconstant x T P independent of TaP but depends on volume ii PVTconstant lll PfoTf PiViTi lV VfVi x Pi Pf x TfTi IV Avogadro s Law Relating Volume and Amount a Law of Combining Volumes volumes of reactant gases at the same pressure and temperature are in ratios of small whole numbers l 2H202 9 2H20 b Avogadro s Law equal volumes of any two gases at the same temperature and pressure contain the same number of molecules i Interpreted from law of combining volumes ii One mole of any gas with the same number of moleculesavogadro s number must occupy the same volume at a given temperature and pressure iii Constant x n V iv HINT cross multiply before you divide with ratios c Molar gas volume Vm volume of one mole of gas 602 x 1023 molecules d Standard Temperature and Pressure STP reference conditions for gases chosen by convention to be 0 C and 1 atm i Typically 224 Lmol or within 2 ii Vm specific constant 224 Lmol STP depending on T and P but independent of gas 53 Ideal Gas Law I Combination of the empirical gas laws ll VmR x TP OR VR x nTP a At one mole the constant R has a specific volume b Temperature must be in Kelvin and units must match i 0 C 27315 K use decimals if C temp given in decimal c Molar Gas Constant R constant of proportionally relating the molar volume of a gas to TP i 082058 LatmKmol most gas law problems ii 83145 JKmol when energy is involved iii 83145 kgm2szKmol iv 83145 kPadm3Kmol v 19872 calKmol d Rewritten as PVnRT ideal gas law e Accurate for low to moderate pressure and low temperature III Calculations Using the Ideal Gas Law a Need three variables P V n or T to calculate the unknown variable IV Gas Density MolecularMass Determination a Gas density varies with temperature and pressure b To obtain molecular mass i Calculate moles in given volume ii Molar mass mass molar volume iii gmol amu or Mmmn C Cr these equations i PVmMmRT ii PMmmvRT iii PMmdRT 54 Stoichiometry Problems Involving Gas Volumes 1 Break it into two problems a Stoichiometry to get the n value b Ideal gas law Perform stoichiometry problem to get number of moles Rearrange ideal gas law for unknown and convert values Plug in n and two other values T V P Insert correct R value Solve P P PWIquot 55 Gas Mixtures Law of Partial Pressures I Experimental conclusion each gas in an unreactive mixture as far as pressure is concerned acts as though it is the only gas in the mixture a Gas at 152 mmHg and gas at 608 mmHg are mixed results in 760 mmHg gas II Partial Pressures and Mole Fractions a Partial Pressure pressure exerted by a particular gas in a mixture b Dalton s Law of Partial Pressures the sum of partial pressures of all the different gases in a mixture is equal to the total pressure of the mixture i PPaPbPc ii Must keep volume the same iii Individual pressure s follow the ideal gas law where ha is the number of moles of that individual gas iv Mole fraction fraction of moles of that component in the total moles of gas mixture 1 Fraction of molecules that are component molecules 2 AnanPaP 3 Mole percent mole fraction x 100 c Collecting Partial Pressure according to OWL 1Calculate the number of moles of each gas from the number of grams and molar mass 2Calculate mole fraction of each gas 3Calculate partial pressure of a gas from its mole fraction and total pressure 4Use Dalton s law for remaining gas II Collecting Gases Over Water a Application of laws of partial pressures b Partial pressure of water vapor in the gas mixture in the collection tube depends on temperature which gives vapor pressure 1Example on page 200 of textbook 56 Kinetic Molecular Theory Kinetic Theory I Kineticmolecular Theory a gas consists of molecules in constant random motion a Kinetic means in motion i Kinetic energy is associated with motion of object of mass m 1 Ek 12 m x speed2 measured in joules speed 12 kgms2 ii Newton s comments pressure due to mutual repulsion of gas molecules which pushed them against sides of containers wrong 1 Not kinetic theory but dominated until the 19th century II Postulates of Kinetic Theory a Physical theories given in terms of postulates i Postulates basic statements from which all conclusions and predictions are deduced b Kinetic Theory of Ideal Gas Five postulates i Most of volume is empty space so volume can usually be ignored ii Molecules move randomly in straight lines in all directions at various speeds Properties that depend on motion ie pressure are the same in all directions iii Intermolecular forces are too weak to matter except in a collision Molecules will keep the same motion until a collision occurs iv Molecular collisions are elastic kinetic energy in a collision remains constantno energy is lost Energy can only be removed by external forces ie heat v Average kinetic energy of a molecule is proportional to absolute temperature The higher the temperature the greater the kinetic energy Establishes temperature from a molecular point of view III Qualitative Interpretation of Gas Laws a Pressure results from bombardment of container walls with molecules kinetic theory i Factors concentration of molecules average speed of molecules determine frequency of collisions 1 Average molecular speed determines force of collision b Relating to Avogadro s Law i More moles increase molecules per unit which in turn increases frequency of collisions with container walls which increases the pressure BUT avogadro s law requires constant pressure so volume increases until concentration decreases ii Result increase in moles at constant pressure and temperature means an increase in volume c Boyle s Law i Increase in volume with constant temperature and number of moles means a decrease in concentration and subsequently a pressure decrease d Charle s Law i Raise the temperature with all other factors fixed which causes an increase in collision force a pressure increase and a volume increase IV Ideal Gas from a Kinetic Theory a According to kinetic theory pressure is proportional to frequency of collision times average force i Average force depends on mass m and average speed u average momentummu 1 The greater the mass and the faster its moving the greater the force 2 An increase in mass an increase in momentum ii Frequency depends on average speed inversely proportional to volume proportional to number of molecule N iii P d u x 1v x Nmu OR PV q Nmu2 1 mu2 is average kinetic energy temperature a PV q NT 2 Number of molecules number of moles a PV q nT 3 Write as an equation by inserting constant a PVnRT 57 Molecular Speeds Diffusion and Effusion I Molecular Speeds a Speeds of molecules vary over a range of values b At any temperature molecular speeds vary widely but most are close to the average which is the maximum in the distribution curve i Increase in temperature means an increase in average speed ii Range of speeds increase as temperatures increase c Rootmeansquare RMS molecular speed u i Type of average molecular speed of a molecule having the average molecular kinetic energy ii U3RTMm12 1 Must be in kg 2 As Mm increases molecular speed decreases 3 Results from postulate 5 4 Use consistent SI units II Diffusion and Effusion a Diffusion process in which a gas spreads out through another gas to occupy a space evenly i Individual molecule moves chaotically bouncing off other molecules and the motion eventually results in complete mixing of the gases ii Depends on 1 Average molecular speed 2 Effect of molecular collisions iii Measured by placing wet indicator strip in air above solution and timing its change as molecules move upwards through air b Effusion process in which a gas flows through a small hole in a container i Graham s Law of Effusion rate of effusion from a particular hole is inversely proportional to the square root of the molecular mass at constant temperature and pressure 1Mm12 ii Three Factors 1 Crosssectional area of hole a Larger hole more molecules escape 2 Number of molecules per unit volume a More crowded more likely to find the hole 3 Average speed a Molecules find the hole sooner iii When comparing bigger mass on bottom 1 Gives how much faster 58 Real Gases I Ideal gas describes behavior of real gases well at low temperatures and moderate temperatures but not high pressures and low temperature Boyle s Law a Deviation s differ for each gas II Applying postulate one high pressure makes volume of molecule important because the space through which a molecule can actually move is very different than the volume III Postulate three intermolecular forces are important at high pressure when the molecules are close together and pressure is less than predicted by ideal IV Van der Waals Equation equation similar to ideal gas law but has constants a and b to account for deviations of real gases from ideal behavior a V become Vnb b P becomes nZav2 P c P nzav2VnbnRT OR P nRTVnb nzav2 i Constants a and b chosen to fit the experiment as closely as possible table on page 210 59 Causes of Gas Becoming a Liquid lecture materialreview I Breakdown of intermolecular forces a Van der Waals corrections II Increase in pressure III Decrease in temperature IV Behaving unideally Chapter Twelve Solutions 121 Types of Solutions I Solution formation a Solution homogenous mixture of two or more substances consisting of ions or molecules i Colloid is similar in appearance II Solute in a solution of a gas or solid dissolved in a liquid is what s dissolved or the thing in a smaller amount III Solvent in a solution of a gas or solid dissolved in a liquid is the liquid or the thing in larger amount IV Gaseous Solutions a Miscible fluids fluids that mix or dissolve in each other in all proportions gases are miscible i mmiscible form two separate layers water and oil b Nonreactive gasesvapors can mix in all proportions to form a mixture V Liquid Solutions a Most obtained from dissolving gas liquid or solid in a liquid i Can also form from combining two solids VI Solid Solutions possible a Usually alloys ie dental fillings 122 Solubility and the Solution Process I Solubility and Saturated Solutions a Dynamic equilibrium of solubility rate at which ions leave crystals rate at which ions return to crystals ie NaCl 6 Na aq Cl39 aq b Saturated solution solution that is in equilibrium with respect to a given dissolved substance Solubility of a solute in water the amount that dissolves in a given quantity of water at a given temperature to give a saturated solution d Supersaturated Solution solution that contains more dissolved substance than a saturated solution does i Not in equilibrium with solid substance ii Crystallization quite fast and dramatic II Factors in Explaining Solubility a Like dissolves like similar substances dissolve one another b Factors i Natural Tendency of Substances to Mix or the Natural Tendency Towards Disorder random movement of molecules ultimately mix them together ii Relative Forces of Attraction between Species lowest energy of solutesolvent obtained 1 Gases have weak intermolecular forces so this doesn t affect them III Molecular Solutions a Gases miscible due to weak intermolecular forces b If the forces are equal then the two move freely about c Similar intermolecular forces means they are soluble in each other d The tendency of molecules to mix results in miscibility IV Ionic Solutions a Differ a lot between their solubility due to different energies of attraction between ions in crystal and ions in water iondipole force i Iondipole force polar water molecules tend to orient with respect to nearby ions and result in an attraction between the ion and the water molecule b Hydration attraction of ions for water molecules favors the dissolving of ionic solid in water i One factor in a hydration of ions is Lattice energy 1 Lattice energy energy holding ions together in the crystal lattice a Works against solution process b Depends on charge of ions and distance between centers of neighboring ions c Inversely proportional to distance between neighboring ions sum of radii of ions d Dominates solubility trends with hydroxides 123 Effects of Temperature and Pressure on Solubility Temperature Change a Most gases become less soluble in water at high temperatures b Ionic substances become more soluble in water as temperature increases i Some decrease though ii Heat sometimes absorbed heat of solution iii Solution becomes hot exothermic v Solution becomes cold endothermic Pressure Change a Little effect on solubility of a liquid or solid but large effect on a gas b Le Chatelier s Principle when a system in equilibrium is disturbed by change of temperature pressure or concentration variable the system shifts in equilibrium composition in a way that tends to counteract this change of variable with a gas c All gases become more soluble in a liquid at a given temperature when the partial pressure of the gas over a solution is increased i Opening a popcan Henry s Law solubility of a gas is directly proportional to the partial pressure of the gas above the solution a SkHP i S is solubility ii kH is Henry s Law constant 1232 Colligative Properties Colligative properties properties that depend on the concentration of only the solute moleculesions in the solutions but not on the chemical identity of the solute a Same structure constant times concentration 124 Ways of Expressing Concentration 0 Concentration amount of solute dissolved in a given quantity of solvent or solution 0 Expressed in volume mass or molar amount 0 Molarity M moles of solute in a liter of solution 0 moles solutesliters solution partvolume 0 Useful for dispersing amount of solute 0 Mass Percentage of Solute 0 Percentage of mass of solute contained in a solution 0 Moles solutemass solution x 100 partwhole o Molality o Moles of solute per kilogram solvent 0 Molal m moles solutekg solvent partpart 0 Independent of temperature 0 Mole fraction 0 the moles of component substance A divided by the total moles of the solution 0 XAmoles of substance Atotal moles of solution I Multiply by 100 for mole percent 0 Conversion of concentration Units 0 Easy to interconvert when expressed in mass or moles of solutesolution pg 503 o Molality requires density 0 Molarity and Molality are almost equal in dilute aqueous solutions 125 Vapor Pressure of a Solution 0 Vaporpressure lowering colligative property equal to the vapor pressure of the pure solvent minus the vapor pressure of the solution 0 If the vapor pressure is lower that means a smaller portion evaporated and less evaporation happens on the surface 0 Raoult s Law Partial pressure of solvent PA over a solution equals the vapor pressure of the solvent P A times the mole fraction of solvent XA in the solution 0 APAP AXA I If nonvolatile PA is total vapor pressure I Observed to hold for dilute solutions I If solvent and solute are similar Raoult s holds true for all mole fractions 0 Assuming solution is nonvolatile nonelectrolyte and Raoult s holds true Vapor pressure lowering is o APP APA 9 APP AXB I Ideal solution of A and B both substances follow raoult s law for all values of mole fractions 0 PP AXAP BXB 0 Can be volatile both have significant vapor pressure 0 Can be used to separate volatile substances from liquid 0 Fractional distillation distilling using a fractionating column 126 Boiling Point Elevation and Freezing Point Depression 0 BoilingPoint Elevation colligative property of a solution equal to the boilingpoint of the solution minus the boilingpoint of the pure solvent 0 Proportional to molal concentration cm of a solution ATbKme Kb is the boilingpoint elevation constant that depends on solvent cm is molality TbT bATb always positive I T b is the boiling point of the pure solvent I Tb is boiling point of the solution 0000 o FreezingPoint Depression colligative property of a solution equal to the freezing point of the pure solvent minus the freezingpoint of the solution 0 Proportional to molal concentration 0 ATfKme o Often used to determine species in solution molal concentration molality and molecular mass 127 Osmosis 0 Only colligative property that is not a change 0 Semipermeable membrane allows smaller molecules to pass through it but not large solvated solute particles of large molecular mass 0 Osmosis phenomenon of solvent flow through a semipermeable membrane to equalize the solute concentrations on both sides of the membrane 0 Osmotic Pressure H colligative property of a solution equal to the pressure that when applied to the solution just stops osmosis 0 Related to molar concentration 0 HMRT I Related to ideal gas law PVnRT and MnV 0 Important in biological processes 0 Intravenous injection corneas etc 0 Reverse Osmosis apply great pressure to reverse osmosis solvent flows into a more concentrated solution from a dilute solution 0 Desalinate remove salts from ocean water to make useable 128 Colligative Properties of Ionic Solutions 0 Total concentration of ions not the concentration of ionic solutions that is important in colligative properties 0 ATfleCm O ATbleCm O nlMRT O APlP AXB I Hill is the number ions resulting from each formula unit 0 DebyeHuckle Theory colligative properties of salts can be explained by assuming the salt is completely ionized but the activities effective concentrations of the ions are less than their actual concentrations as a result of electrical interactions between ions 129 Colloids o Colloid dispersion of particles of one substance the dispersed phase throughout another substance or solution the continuous phase 0 ie fog o Differs from a true solution because dispersed particles are larger than normal molecules but too small to be seen with a microscope o Tyndall Effect 0 Colloid has the ability to scatter light unlike a normal solution 0 Tyndall effect scattering of light by colloidsized particles 0 Types of Colloids o Colloids are characterized according to state of dispersed phase and of the continuous phase 0 Aerosols liquid droplets or solid particles dispersed throughout a gas 0 Emulsion liquid droplets dispersed throughout another liquid 0 Sol solid particles dispersed in liquid 0 Hydrophobic and Hydrophilic Colloids o Colloids in continuous phase of water are either hydrophilic or hydrophobic o Hydrophilic colloid colloid in which there is a strong attraction between the dispersed phase and the continuous phase water 0 Hydrophobic colloid colloid in which there is a lack of attraction between the dispersed phase and continuous phase water 0 Coagulation o Coagulation process process by which the dispersed phase of a colloid is made to aggregate and thereby separate from the continuous phase I Positive colloidal ion gathers layer of anions around it 0 Thickness depends on charge of anions 0 Greater the magnitude of negative charge the more compact the layer o If close enough it become neutralized I Can approach close enough to aggregate 0 Association Colloidals o Micelle colloidalsized particle formed in water by the association of molecules or ions that each have a hydrophobic and hydrophilic end I Hydrophobic end faces in and hydrophilic faces out towards water 0 Association colloidal colloid in which the dispersed phase consists of micelles
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