Chapter 4 Solutions _ Their Properties

CHAPTER 4 SOLUTIONS & THEIR PROPERTIES

SOLUTIONS Learning objectives:  Discuss different kinds of solutions

Types of Homogeneous Mixtures Solutions • • •

most important class of homogeneous mixtures contain particles with diameters in range 0.1-2 nm transparent, do not separate on standing (salt water, sugar water)

Colloids • •

contain particles with diameters in the range 2-500nm murky, do not separate on standing (milk, fog,)

Suspensions • • •

having larger particles than colloids Not truly homogeneous Particles separate on standing (blood, paint)

Solutes and Solvents • A solution consists of a solute and a solvent:  solute  the substance which is being dissolved.  solvent  the substance (usually a liquid) that dissolves the solute (usually, the solvent is the most abundant component in the mixture). • Aqueous solution are solutions in which the solvent is water.

Solute

Solvent

Solution

Kinds of Solutions Solution Phase Gaseous solution Liquid solutions

Solid solutions

Solute Solvent

Examples

Gas

Gas

Air (O2, N2, Ar, CO2, H2O, and other gases)

Gas

Liquid

Carbonated water (CO2 in water)

Liquid

Liquid

Gasoline (Mixture of hydrocarbons), vodka (ethanol and water)

Solid

Liquid

Seawater (NaCl and other salts in water)

Gas

Solid

H2 in palladium metal

Liquid

Solid

Dental amalgam (mercury in silver)

Solid

Solid

Metal alloys such as sterling silver (Ag and Cu), brass (Cu and Zn) and bronze (Cu and Sn); waxes

ENERGY CHANGES & THE SOLUTION PROCESS Learning objectives:  Identify intermolecular forces in solution  Explain types of solution interactions  Identify intermolecular forces in solutions  Describe dissolution of NaCl in water  Explain the rule of thumb "like dissolves like.“  Define the entropy and enthalpy of solution

Intermolecular Forces in Solutions Relative strengths of intermolecular forces must be considered between solute and solvent particles that promote or prevent the formation of a solution. Intermolecular Forces (in order of decreasing strength):  ion-dipole forces — solvent molecules cluster around ions in hydration shells, disrupting the bonding in the crystal lattice.  Hydrogen bonds— substances with O—H and N—H bonds are often soluble in water because of H-bonding (unless the molecules are large).

Intermolecular Forces in Solutions  dipole-dipole forces —polar solutes interact well with polar solvents through attraction of partial charges.  ion - induced dipole forces — responsible for the attraction between Fe2+ and O2 molecules in the bloodstream.  dipole - induced dipole forces —responsible for the solvation of gases (nonpolar) in water (polar).  London (dispersion) forces —the principal attractive force in solutions of nonpolar substances (e.g., petroleum).

Solution Interactions

Solutions form when solvent-solvent, solute-solute, and solute-solvent forces are similar.

Solution Interactions Relative Interaction and Solution Formation Solvent-solute interaction > Solvent-solvent and solute-solute interaction Solvent-solute interaction = Solvent-solvent and solute-solute interaction Solvent-solute interaction < Solvent-solvent and solute-solute interaction

Solution forms

Solution forms

Solution may or may not form depending on relative disparity

Dissolution of NaCl in water

The General Solubility Rule The general rule in solubility is that “like dissolves like” Water, a polar molecule, dissolves ethanol, which is also polar, but does not dissolve hexane and dichloromethane which are both nonpolar.  Ethanol and water are miscible—completely soluble in each other in all proportions.  Hexane and water are immiscible—they do not mix with each other at all.  Hexane and dichloromethane are miscible with each other.

Entropies of solutions Entropy is a measure of the disorder or energy randomization in a system. Entropies of solution are usually positive because molecular randomness usually increases when (a) a solid dissolves in a liquid or (b) one liquid dissolves in another.

Enthalpy of solution •

Enthalpy of solution measures how much energy is either absorbed or released when a solution is prepared.

•

The value of ΔHsoln is the sum of three terms: Solvent-solvent interactions: Energy is required (+ ΔH) to overcome intermolecular forces between solvent molecules because the molecules must be separated and pushed apart to make room for solute particles Solute-solute interactions: Energy is required (+ ΔH) to overcome intermolecular forces holding solute particles together in a crystal. Solvent-solute: Energy is released (- ΔH) when solvent molecules cluster around solute particles and solvate them

Energy Changes and the Solution Process The solute-solvent interactions are greater than the sum of the solute-solute and solventsolvent interactions.

Energy Changes and the Solution Process The solute-solvent interactions are less than the sum of the solute-solute and solvent-solvent interactions.

Enthalpy of solution

UNITS OF CONCENTRATION Learning objectives:  Interconvert units of concentration  Perform calculations using solution density, molarity,mole fraction, weight percent, parts per million, parts per billion, and molality.

Units of Concentration • Concentration: The amount of solute present in a given amount of solution. • Molarity (M): Molarity 

Moles of solute Liters of SOLUTION

• Mole Fraction (X):

XA 

Moles of A Total number of moles

Units of Concentration • Mass percent: The ratio of the mass of a solute to the mass of a solution, multiplied by 100%.

mass of solute % by massof solute = 100%  mass of solution mass of solution =mass of solute +mass of solvent

Units of Concentration • Parts per Million (ppm):

Mass of solute x 10 6 Total mass of solution

• One ppm gives 1 gram of solute per 1,000,000 g or one mg per kg of solution. For solid samples: ppm = µg/g = mg/kg • For dilute aqueous solutions this is about 1 mg per liter of solution. For liquid samples: ppm = µg/mL = mg/L

Units of Concentration • Parts per Billion (ppb):

For solid samples:

Mass of solute x 109 Total mass of solution

ppm = µg/g = mg/kg

For liquid samples: ppm = µg/mL = mg/L • Molality (m): Moles of solute Molality = Kilograms of SOLVENT

Units of Concentration

FACTORS AFFECTING SOLUBILITY Learning objectives:  Define saturated, unsaturated and supersaturated solutions  Describe crystallization process  Discuss the effect of temperature and pressure on solubility  State Henry’s Law and its examples

Saturated and Unsaturated Solutions • Saturated: Contains the maximum amount of solute that will dissolve in a given solvent. • Unsaturated: Contains less solute than a solvent has the capacity to dissolve. • Supersaturated: Contains more solute than would be present in a saturated solution; these solutions are unstable, and a slight disturbance causes the “extra” solute to precipitate out. • Crystallization: The process in which dissolved solute comes out of the solution and forms crystals.

Precipitation from a supersaturated solution a)

A supersaturated solution of sodium acetate in water

b)

When a tiny seed crystal is added, larger crystals begin to grow and precipitate from the solution until equilibrium is reached

Effect of Temperature on Solubility Solubilities are temperature-dependent. The solubility of most molecular and ionic solids increases with temperature, although some are almost unchanged, and some decrease For a solute with ΔHsoln > 0: solute + solvent + heat saturated solution  solubility increases with temperature. For a solute with ΔHsoln < 0: solute + solvent saturated solution + heat  solubility decreases with temperature.

Effect of Temperature on Solubility

Solids:

Solubilities of some common solids in water as a function of temperature. Most substances become more soluble as temperature rises, although the exact relationship is often complex and nonlinear.

Effect of Temperature on Solubility

Gases:

• Solubilities of some gases in water as a function of temperature. • Most gases become less soluble in water as the temperature rises.  Soft drinks become “flat” as they warm up and lose carbon dioxide.  Aquatic life is affected by decreasing amounts of dissolved oxygen as a result of thermal pollution.

Effect of Pressure on the Solubility Pressure has little effect on the solubility of solids and liquids, but has a large effect on gases. At a given pressure, there is an equilibrium between the gas which is dissolved in the solution and the gas in the vapor phase. If the pressure increases, more gas dissolves to reduce the “extra” pressure; the new equilibrium is established with more gas dissolved.

Effect of Pressure on the Solubility • Henry’s Law: The solubility(Sgas, in mol/L) of a gas is proportional to the pressure of the gas (Pgas, in atm) over the solution Sgas = k . Pgas (The Henry’s law constant, k is a proportionality constant, unique to each gas, at a given temperature, with units of mol L-1 atm-1.)

Effect of Pressure on the Solubility Examples of Henry’s – law behavior: • When a can of soda is opened, bubbles of gas fizzing out of solution because the pressure of CO2 in the can drops and CO2 suddenly becomes less soluble. • If a deep sea diver comes up to the surface too quickly, N2 which has dissolved in his bloodstream at higher pressures comes back out of solution.  The N2 forms bubbles which block capillaries and inhibit blood flow, resulting in a painful, and potentially lethal, condition called the “bends.”  Less soluble gases, such as He, are often used in the breathing mixtures to reduce this problem.