Solutions Chapter 12.

Name: Solutions Chapter 12.
Uploaded: 2017-07-14T15:14:05+00:00
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Description: Solutions Chapter 12.

Solutions Chapter 12

Classification of Matter
Solutions are homogeneous mixtures

12.1 Solution Concentration
Objectives To express the concentration of a solution in several different ways To convert between different concentration units

The composition of a solution is expressed as concentration Concentration is always expressed as the quantity of solute present in a fixed quantity of the solution or solvent Remember your definitions from Ch 4!

Solute Solvent A solute is the dissolved substance in a solution.
Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water in soda

Types of Solutions

Although molarity (moles of solute per liter of solution) is the preferred concentration unit for stoichiometry calculations, other concentration units are important for other types of calculations Ex: use of molality in determining changes in boiling point and freezing point since it is not affected by changes in temperature as is molarity The following slides illustrate the numerous ways to denote concentration……

Calculations of Solution Concentration
Molarity - the ratio of moles of solute to liters of solution

Calculate the molarity of KOH in a solution prepared by dissolving 8.23 g of KOH in enough water to form 250 mL of solution. A: M Calculate the number of moles of solute in 33 mL of a 3.11 M HNO3 solution. A: mol

Mole fraction – the ratio of moles of solute to total moles of solution

Mole Fraction, cont’d ΧA + ΧB + ΧC + ….. = 1
Please note that more than two substances can be involved and that the sum of all mole fractions equals 1 ΧA + ΧB + ΧC + ….. = 1 Note we used these types of calculations when we calculated partial pressures of one component in a gas mixture

What is the mole fraction of each gas in a mixture that contains 2.3 g of neon, 0.33 g of xenon, and 1.1 g of argon? A: Moles of each are as follows: Ne: moles Xe: moles and Ar: moles Therefore mole fractions are as follows: Ne: Xe: Ar: 0.19

Mass percent - the ratio of mass units of solute to mass units of solution, expressed as a percent

Determine the mass percent of a solution prepared by dissolving 5.00 g of NaCl in 200 g of water. A: 2.44% A solution of a mass of 25.0 g contains 2.00 g of glucose. Express the concentration in mass percent of glucose. A: 8.00%

Molality – moles of solute per kilogram of solvent

Be careful not to confuse molarity (M) and molality (m) Determine the molality of the solution made with 5.00 g of NaCl in 200 g of water. A: m What mass of acetic acid (CH3CO2H, molar mass = g/mol) must be dissolved in 250 g of water to produce m solution? A: g

How many grams of water must be added to 4.00 g of urea (CO(NH2)2); molar mass = g/mol) to produce a m solution of the compound? A: 266 g

Need to be able to convert from one concentration designation to another Remember to write the given concentration as a fraction and the desired units as a fraction and figure out how to get from point A to point B

Find the (a) molality and (b) mole fraction of a 24.5% solution of ammonia (NH3) in water A: (a) m (b) 0.256 Find the molarity of a 3.10 molal aqueous ammonia solution that has a density of g/mL. A: M

12.2 Principles of Solubility
Objectives To predict the relative solubilities of substances in different solvents based on solute-solvent interactions To explain how disorder is a driving force in the mixing of substances

For most solutes, there is a limit to the quantity that can dissolve in a fixed volume of a given solvent. The solubility of a solute is reached when the rate of dissolution and the rate of crystallization are equal

Saturation of Solutions
A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated. A solution that contains less solute than a saturated solution under existing conditions is unsaturated. A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.

The addition of a small quantity of solute to a solution is a simple way to distinguish among unsaturated, saturated, and supersaturated solutions

Crashing Out of a Supersaturated Solution

Solutions play a major role in chemical reactions Two of the most important factors that determine whether a given substance will dissolve in a solvent include: the enthalpy change that accompanies solute-solvent interactions and the change in disorder

Steps in Solution Formation
H1 Expanding the solute (requires input of energy) Separating the solute into individual components H2 Expanding the solvent (requires input of energy) Overcoming intermolecular forces of the solvent molecules H3 Interaction of solute and solvent to form the solution (releases energy) And: ΔHsoln = ΔH1 + ΔH2 + ΔH3

Fig. 12-5, p. 477

Heat of Solution The Heat of Solution is the amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent. Substance Heat of Solution (kJ/mol) NaOH -44.51 NH4NO3 +25.69 KNO3 +34.89 HCl -74.84

Surprisingly many substances with a + ΔHsoln still dissolve spontaneously. This is due to an increase in entropy (disorder) which is favorable

Solubility of Molecular Compounds When trying to predict solubility you need to ask several questions… What intermolecular forces need to be broken? What intermolecular forces are being generated? Are the entropy increases sufficient to overcome any unfavorable ΔH values?

Predicting Solubilities…
Predict the solvent of each pair in which the given compound is more soluble CCl4 in water or hexane (C6H14) Urea (CO(NH2)2) in water or CCl4 Iodine (I2) in benzene or water

Solubility of Ionic Compounds in Water Note that water surrounds ions in a sphere of hydration. Note that the Hδ+ of water molecules are oriented towards the anions in solution and the Oδ- are oriented around the cations. Note that disorder increases significantly when ionic solids dissolve in solution.

Predicting Solution Formation
Solvent/ Solute H1 solute H2 solvent H3 Hsol’n Outcome Polar/ Polar + large - large +/-small Solution forms Nonpolar + small +/- small No solution Nonpolar/ +/- small polar

12.3 Effects of Pressure and Temperature on Solubility
Objectives To state the effects of pressure and temperature on solubility To calculate the solubility of gases using Henry’s Law To explain why the solubilities of solids and liquids do not change appreciably with changes in pressure To relate the sign of the enthalpy of solution to the increase or decrease in solubility with temperature

Effect of Pressure on Solubility Pressure has very little effect on the solubilities of solids and liquids The solubility of a gas is directly proportional to its partial pressure at a given temperature / [A] is calculated via Henry’s law (CA = kPA) k = a proportionality constant that is characteristic of the particular solute, solvent, and temperature For gaseous solutes, an increase in pressure is relieved by additional gas dissolving in the liquid

Henry’s Law Constants

What is the molal concentration of oxygen in water at 20°C that has been saturated with air at 1 atm? Assume that the mole fraction of oxygen in air is 0.21. Equation: C = kP k for oxygen at 20°C is 1.43 x 10-3 molal/atm P = mole fraction x total pressure = 0.21 x 1 atm = 0.21 atm C = 1.43 x 10-3 molal/atm • 0.21 atm = 3.0 x 10-4 molal

What is the molal concentration of nitrogen in water at 20°C that has been saturated with air at 1 atm? Assume that the mole fraction of nitrogen in air is 0.79. Equation: C = kP k for nitrogen at 20°C is 7.34 x 10-4 molal/atm P = mole fraction x total pressure = 0.79 x 1 atm = 0.79 atm C = 7.34 x 10-4 molal/atm • 0.79 atm = 5.8 x 10-4 molal

Predicting the solubility of solids or gases with changes in temperature
To predict solubilities, I find it helpful to think of heat as a reactant (endothermic reactions) or a product (exothermic reactions) Le Chatelier’s Principle can then be used to predict what is favored A + solvent + heat ↔ solution B + solvent ↔ solution + heat

As temperature increases, the solubility increases for any substance with an endothermic enthalpy of solution and decreases for one with an exothermic enthalpy of solution. A + solvent + heat ↔ solution B + solvent ↔ solution + heat

A + solvent + heat ↔ solution B + solvent ↔ solution + heat Since the enthalpy of solution for most gases in water is exothermic their solubilities decrease with an increase in temperature. The solubilities of most solids increases as the temperature of the solution increases. Some substances that appear to violate these generalizations undergo a more complex dissolution process.

Solubility Chart

The enthalpy of solution of potassium chromate (K2CrO4) in water is kJ/mol. How does the solubility of this compound change when the temperature is lowered? Decrease At 1 atm pressure, the solubility of oxygen is x 10-3 molal at 20°C and 8.71 x 10-4 molal at 60°C. What is the sign of the enthalpy of solution of oxygen? Negative/Exothermic

Solubility Trends The solubility of MOST solids increases with temperature. The rate at which solids dissolve increases with increasing surface area of the solid. The solubility of gases decreases with increases in temperature. The solubility of gases increases with the pressure above the solution.

Therefore… Solids tend to dissolve best when: Heated Stirred
Ground into small particles Gases tend to dissolve best when: The solution is cold Pressure is high

12.4 Colligative Properties of Solutions
Objectives To define and identify the colligative properties of solutions To relate the values of colligative properties to the concentration of solutions To calculate the molar masses of solutes from measurements of colligative properties

Colligative Properties of Solutions
Some questions to think about…. Why do we use CaCl2 in the winter time? Does it take longer for salted water or pure water to boil?

Colligative Properties are those properties of solution that change in proportion to the concentration of solute particles and are not based upon the identity of those particles. Colligative properties include: Vapor Pressure Lowering Boiling Point Elevation Freezing Point Depression Osmotic Pressure

All of the colligative properties fit the relationship: property = solute concentration x constant They differ in the units in which the solute concentration is expressed

Vapor Pressure of the Solvent Adding a nonvolatile solute to a solvent always reduces the equilibrium vapor pressure below that of the pure solvent

Effect of solute on vapor pressure
Pure Solvent Solution

Vapor Pressure of the Solvent Raoult’s law expresses the relationship between vapor pressure of a pure solvent and the vapor pressure of a solution

Raoult’s Law Psolvent = Observed Vapor pressure of
The presence of a nonvolatile solute lowers the vapor pressure of the solvent. Psolvent = Observed Vapor pressure of the solution solvent = Mole fraction of the solvent P0solvent = Vapor pressure of the pure solvent

mole fraction solute in
The relationship between vapor pressure of a pure solvent and the vapor pressure of a solution Psolvent = Xsolvent x P °solvent a = pure benzene mole fraction solute in b = 0.02 c = 0.04 d = 0.06 e = 0.08

If the term Xsolvent is replaced by (1-Xsolute) and we rearrange the equation: ΔP = P°solvent – Psolvent = Xsolute•P°solvent

At 25°C the vapor pressure of pure benzene is 93.9 torr. A solution of a nonvolatile solute in benzene has a vapor pressure of 91.5 torr at the same temperature. What is the concentration of the solute expressed as a mole fraction? ΔP = Xsolute•P°solvent Xsolute = 0.026 What is the solute concentration in a benzene solution that has a vapor pressure of 90.6 torr at 25°C? Xsolute = 0.035

Boiling Point Elevation Caused by the fact that solute particles “hold onto” solvent molecules and make it less likely for liquid solvent molecules to move to the gas phase Equation: ΔTb = mkb Where ΔTb = elevation of the boiling point, m = the molality of the solute, and kbis the boiling point elevation constant depending on the solvent

Boiling Point Elevation of Water
Each mole of solute particles raises the boiling point of 1 kilogram of water by 0.51 degrees Celsius. kb = 0.51 C  kilogram/mol m = molality of the solution i = van’t Hoff factor = 1 if a non-electrolyte is the solute

A solution is prepared by dissolving 1.00 g of a nonvolatile solute in 15.0 g of acetic acid. The boiling point of this solution is °C. The normal boiling point of acetic acid is °C and kb is 3.07°C/m. Find the molar mass of the solute. A: molar mass = 90.1g/mol Find the molar mass of a nonvolatile solute, if a solution of 1.20 g of the compound dissolved in 20.0 g of benzene has a boiling point of 80.94C. The normal boiling point of benzene acid is 80.10°C and kb is 2.53°C/m. A: molar mass = 1.8 x 102 g/mol

Freezing Point Depression When a solute is present, both the vapor pressure of the solvent and the triple point temperature is lowered. As a result, freezing point is also depressed. Temperature °C Pressure, atm 1 atm

Equation: ΔTf = mkf Where ΔTf = depression of the freezing point, m = the molality of the solute, and kf is the freezing point depression constant depending on the solvent

BP Elevation & FP Depression Constants

Freezing Point Depression of Water
Each mole of solute particles lowers the freezing point of 1 kilogram of water by 1.86 degrees Celsius. kf = 1.86 C  kilogram/mol m = molality of the solution i = van’t Hoff factor = 1 if the solute is a non-electrolyte

Pure ethylene dibromide freezes at 9.80°C. A solution made by dissolving g of ferrocene (molecular formula Fe(C5H5)2, molar mass = g/mol) in 10.0 g of ethylene dibromide has a freezing point of 8.45°C. What is the freezing point depression constant for the solvent ethylene dibromide? A: kf = 11.8°C/m Benzophenone has a freezing point of 49.00°C. A molal solution of urea in this solvent has a freezing point of 44.59°C. Find the freezing point depression constant for the solvent benzophenone. A: kf = 9.80°C/m

Osmotic Pressure Semi-permeable membranes allow some substances to move across but not others and are especially important to transport in biological systems Osmosis = the diffusion of solvent across a semi-permeable membrane Osmotic pressure = the pressure difference needed for no net transport of solvent to occur across a semi-permeable membrane that separates the solution from the pure solvent

Osmotic Pressure H2O

Osmotic Pressure Calculations
M = Molarity of the solution (= n/V) R = Gas Constant = Latm/molK i = van’t Hoff factor = 1 for non-electrolytes

Of all the colligative properties, osmotic pressure is the most sensitive. Because it is so sensitive it used to measure the molar masses of very large molecules and substances that are only slightly soluble in water.

Hemoglobin is a large molecule that carries oxygen in human blood. A water solution that contains g of hemoglobin in 10.0 mL of solution has an osmotic pressure of 7.51 torr at 25°C. What is the molar mass of hemoglobin? A: molar mass = 6.51 x 104g/mol A 5.70 mg sample of a protein is dissolved in water to give 1.00 mL of solution. If the osmotic pressure of this solution is 6.52 torr at 20°C, what is the molar mass of the protein? A: molar mass = 1.60 x 104g/mol

12.5 Colligative Properties of Electrolyte Solutions
Objectives To predict the ideal van’t Hoff factor of ionic solutes To calculate the expected colligative properties for solutions of electrolytes

Solutions of ionic and molecular solutes behave differently. van’t Hoff ( ) noticed that the effect of some solutes on colligative properties was greater than expected van’t Hoff factor = i Works best for dilute electrolyte solutions = 1 for nonelectrolytes Ideal value = # of particles formed when electrolytes dissociate Ideal value of i is typically greater than the measured value of i

Dissociation Equations and the Determination of i
NaCl(s)  Na+(aq) + Cl-(aq) i = 2 AgNO3(s)  Ag+(aq) + NO3-(aq) i = 3 MgCl2(s)  Mg2+(aq) + 2 Cl-(aq) i = 3 Na2SO4(s)  2 Na+(aq) + SO42-(aq) AlCl3(s)  Al3+(aq) + 3 Cl-(aq) i = 4

Ideal vs. Real van’t Hoff Factor
The ideal van’t Hoff Factor is only achieved in VERY DILUTE solution.

Arrange the following solutions in order of increasing osmotic pressure: M sucrose, 0.02 M HNO3, 0.01 M BaCl2 A: 0.02 M sucrose, 0.01 M BaCl2 , 0.02 M HNO3

12.6 Mixtures of Volatile Substances
Objectives To calculate the vapor pressure of each component and the total vapor pressure over an ideal solution To interpret positive and negative deviations from Raoult’s law in terms of the relative strengths of the intermolecular attractions in the solution

At times, liquid solutions can contain two or more volatile components. Raoult’s Law can be modified to account for more than one volatile component in a solution The vapor over a solution of two volatile components contains a larger fraction of the more volatile substance than does the solution.

Liquid-liquid solutions in which both components are volatile
Modified Raoult's Law: P0 is the vapor pressure of the pure solvent PA and PB are the partial pressures

Vapor Pressure of an Ideal Solution
Total Pressure Benzene Pressure (torr) Toluene Mole fraction of toluene

At 60°C the vapor pressure of benzene is 384 torr and that of toluene is 133 torr. A mixture is made by combining 1.20 mol of toluene with 3.60 mol of benzene. Find: (a) The mole fraction of toluene in the liquid (b) The partial pressure of toluene above the liquid (c) The partial pressure of benzene above the liquid (d) The total vapor pressure (e) The mole fraction of toluene in the vapor phase

It is important to note that not all volatile components behave ideally throughout the range of composition (mole fraction = 0 to mole fraction = 1) Positive deviations indicate the observed vapor pressure is greater than expected. Occurs when A-B interactions are weaker than the average of the attractions in the pure components of the mixture Negative deviations indicate the observed vapor pressure is less than expected Occurs when intermolecular forces between dissimilar molecules are stronger than the average intermolecular forces in the pure substances

Positive Deviations from Raoult’s Law
Weak solute-solvent interaction results in a vapor pressure higher than predicted Endothermic mixing = Positive deviation

Negative Deviations from Raoult’s Law
Strong solute-solvent interaction results in a vapor pressure lower than predicted Exothermic mixing = Negative deviation

Fractional distillation allows two volatile liquids to be separated. Enrich the vapor phase with more volatile compound during each successive evaporation/conden-sation step

Azeotrope Many liquids form constant boiling mixtures called azeotropes The composition of the liquid and vapor phases are the same Complete separation of volatile liquids that form an azeotrope cannot be achieved by fractional distillation

Example of an azeotrope: HNO3 Pressure 1 atm 0.4 1.0 Mole fraction of HNO3

Additional slides follow that review some other terms/examples related to solutions but are not presented in Reger’s Ch 12….

Definition of Electrolytes and Nonelectrolytes
An electrolyte is: A substance whose aqueous solution conducts an electric current. A nonelectrolyte is: A substance whose aqueous solution does not conduct an electric current. Try to classify the following substances as electrolytes or nonelectrolytes…

Electrolytes? Pure water Tap water Sugar solution
Sodium chloride solution Hydrochloric acid solution Lactic acid solution Ethyl alcohol solution Pure sodium chloride

Answers to Electrolytes
NONELECTROLYTES: Tap water (weak) NaCl solution HCl solution Lactate solution (weak) Pure water Sugar solution Ethanol solution Pure NaCl

Electrolytes vs. Nonelectrolytes
The ammeter measures the flow of electrons (current) through the circuit. If the ammeter measures a current, and the bulb glows, then the solution conducts. If the ammeter fails to measure a current, and the bulb does not glow, the solution is non-conducting.

Osmotic Pressure The minimum pressure that stops the osmosis is equal to the osmotic pressure of the solution

Suspensions and Colloids
Suspensions and colloids are NOT solutions. Suspensions: The particles are so large that they settle out of the solvent if not constantly stirred. Colloids: The particles intermediate in size between those of a suspension and those of a solution.

Types of Colloids Examples Dispersing Medium Dispersed Substance
Colloid Type Fog, aerosol sprays Gas Liquid Aerosol Smoke, airborn germs Solid Whipped cream, soap suds Foam Milk, mayonnaise Emulsion Paint, clays, gelatin Sol Marshmallow, Styrofoam Solid Foam Butter, cheese Solid Emulsion Ruby glass Solid sol

The Tyndall Effect Colloids scatter light, making a beam visible. Solutions do not scatter light. Which glass contains a colloid? colloid solution

Solutions Chapter 12.

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