Mixtures and Solutions Section 14-1 Mixtures Heterogeneous vs. HomogeneousHeterogeneousHomogeneous A heterogeneous mixture is a mixture that does not.

Mixtures and Solutions

Mixtures

• Heterogeneous vs. Homogeneous

• A heterogeneous mixture is a mixture that does not have a uniform composition and in which the individual substances remain distinct.

• Suspensions are heterogeneous mixtures containing particles that settle out if left undisturbed.

• Examples: sand and silt in river water, paint, etc.

Heterogeneous Mixtures (cont.)

• Colloids are heterogeneous mixtures of intermediate sized particles (between 1 nm and 1000 nm) and do not settle out.

• Consist of “dispersed particles” in the “dispersion medium” (the most abundant).

• Milk is a great example of a colloid. Tiny liquid butterfat globules stay suspended in a water dispersion medium.

Heterogeneous Mixtures (cont.)

• Why don’t colloidal particles settle out?

• Brownian motion is the jerky, random movements of particles in a liquid colloid, from the results of particle collisions.

• The Tyndall effect is when dispersed colloid particles scatter light (does not occur with solutions).

• Polar groups on their surface attract polar dispersion medium molecules or dissolved ions.

• Charged layers around particles repel each other and keep particles suspended.

Heterogeneous Mixtures (cont.)

Homogeneous Mixtures

• Solutions are homogeneous mixtures that contain two or more substances called the solute and solvent.

• Term is typically used with liquids (water), but gases and solids form solutions too.

• The solvent is the most abundant material in the mixture. Water is our key solvent.

• Solutes are the less abundant material that is mixed (“dissolved”) in the solvent. • Can be liquids, solids or gases…all dissolve!

Homogeneous Mixtures (cont.)

Solutions

• A substance that dissolves in a solvent is soluble.

• A substance that does not dissolve in a solvent is insoluble.

• Two liquids that are soluble in each other in any proportion are miscible.

• Two liquids that can be mixed but separate shortly after are immiscible.

• Electrolytes– Compounds that dissociate into separate ions

in water and are good conductors.– “Strong” electrolytes = 100% dissociation

• Ionic compounds are strong electrolytes• Also includes the strong acids: HCl, HNO3, HBr

• Fully dissociate to form H+ ions in solution (H3O+)

– “Weak” electrolytes < 100% dissociation• Nonelectrolytes

– Do not dissociate into ions in water. – Nonpolar molecules

Solutions

A. A

B. B

C. C

D. D

Section Assessment

The jerky, random movement of particles in a liquid colloid is known as ____.

A. Brownian motion

B. Tyndall effect

C. Charles’s Law

D. kinetic energy

Expressing Concentration

• The concentration of a solution is a measure of how much solute is dissolved in a specific amount of solvent or solution.

• The level of concentration can be described as concentrated or dilute.

• These are relative (qualitative) terms.

• Quantitative levels of concentration can be described using a variety of ratios.

Expressing Concentration (cont.)

Expressing Concentration (cont.)

Expressing Concentration (cont.)

• Molarity is the number of moles of solute dissolved per liter of solution.

• Dilution equation: M1V1 = M2V2

Expressing Concentration (cont.)

• Molality is the ratio of moles of solute dissolved in 1 kg of solvent.

Expressing Concentration (cont.)

• Mole fraction is the ratio of the number of moles of solute in solution to the total number of moles of solute and solvent.

where XA and XB represent mole fractions of each substance

The Solvation Process

• Solvation is the process of surrounding solute particles with solvent particles to form a solution.

• Solvation in water is called hydration.

• Solvation of Ionic Compounds:

• The attraction between dipoles of a water molecule and the ions of a crystal are greater than the attraction among ions of a crystal.

• Ion-dipole intermolecular force

The Solvation Process (cont.)

The Solvation of Molecules

• Sucrose molecules have several O–H groups, which become sites for hydrogen bonding with water molecules.

• Oil does not form a solution with water because there is little attraction between polar water molecules and nonpolar oil molecules.

The Solvation Process (cont.)

• During solvation, the solute must separate into particles and move apart. The solvent must also move apart to make room for solute particles. Both of these actions require energy (endothermic).

• Then, when the solute and solvent mix and attract each other, energy is released (exothermic).

• The overall energy change that occurs during solution formation is called the heat of solution.

Factors That Affect Solvation

• Agitation: Stirring or shaking moves dissolved particles away from the contact surfaces more quickly and allows new collisions to occur.

• Breaking the solute into small pieces increases surface area and allows more collisions to occur.

• As temperature increases, rate of solvation increases.

Solubility

• Solubility = the maximum amount a solute will dissolve in a solvent. It depends on the type of solute and solvent and temperature.

• As concentration of solute in solvent increases, more solute particles collide with remaining crystalline solid and precipitate.

• Unsaturated solutions are solutions that contain less dissolved solute for a given temperature and pressure than a fully saturated solution.

Solubility (cont.)

• Saturated solutions contain the maximum amount of dissolved solute for a given amount of solute at a specific temperature and pressure. Equilibrium is reached between solvation and precipitation.

• Solubility is affected by increasing the temperature of the solvent because the kinetic energy of the particles increases.

Solubility Graph (Water is the Solvent)

Temperature can make a big difference!

Solubility (cont.)

• A supersaturated solution contains more dissolved solute than a saturated solution at the same temperature.

• To form a supersaturated solution, a saturated solution is formed at high temperature and then slowly cooled. Adding more solid will spark rapid precipitation (‘seeding’).

• Supersaturated solutions are unstable.

Solubility of Gases• Gases are less soluble in liquid solvents at

high temperatures. Higher K.E. allows more gas molecules to escape from solution.

• Solubility of gases increases as its external pressure is increased. Carbonated sodas!

• Henry’s law states that at a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P).

A. A

B. B

C. C

D. D

Section Assessment

For a given amount, which type of solution contains the LEAST amount of solute?

A. solvated

B. saturated

C. supersaturated

D. unsaturated

A. A

B. B

C. C

D. D

Section Assessment

At a given temperature, the solubility of a gas is directly proportional to what?

A. volume

B. mass

C. molarity

D. pressure

Colligative Properties of Solutions

• Colligative properties are physical properties of solutions that are affected by the number of particles but not by the identity of dissolved solute particles.

• Types of colligative properties:• Vapor pressure lowering• Boiling point elevation• Freezing point depression• Osmotic pressure

Colligative Properties of Solutions

• Strong electrolytes have a greater colligative effect since they put more particles into solution than weak electrolytes or molecules.

• The relationship between moles of solute and moles of particles in solution is the van’t Hoff factor:

“ i ” = (moles of particles) / (moles of solute)

• Molecules always have an i = 1

• Strong electrolytes have i = 2 or more depending upon the number of ions in formula (CaCl2 has i = 3).

• Assume all electrolytes are “strong” (i.e. 100% dissociate) unless told otherwise.

Vapor Pressure Lowering

• Adding a nonvolatile solute to a solvent lowers the solvent’s vapor pressure.

• Volatility is tendency of a substance to vaporize (change from liquid to gas).

• When a solute is present, a mixture of solvent and solute occupies the surface area, and fewer particles enter the gaseous state.

• The greater the number of solute particles, the lower the vapor pressure.

Vapor Pressure Lowering (cont.)

• Vapor pressure lowering is due to the number of solute particles in solution.

Vapor Pressure Lowering (cont.)

• Raoult’s Law:

P(solution) = X(solvent) x Po (solvent)

where:

X(solvent) = mole fraction of the solvent

Po (solvent) = vapor pressure of pure solvent

Boiling Point Elevation

• When a nonvolatile solute lowers the vapor pressure of a solvent, the boiling point is also affected.

• Recall that the boiling point is when the vapor pressure of the liquid is equal to the atmospheric pressure.

• More heat is needed to supply additional kinetic energy to raise the vapor pressure to atmospheric pressure.

Boiling Point Elevation (cont.)

• The temperature difference between a solution’s boiling point and a pure solvent's boiling point is called the boiling point elevation.

ΔTb = iKbm

• where ΔTb is the boiling point elevation, i is the van’t Hoff factor, Kb is the molal boiling point elevation constant, and m represents molality.

Boiling Point Elevation (cont.)

Freezing Point Depression

• At a solvent's freezing point temperature, particles no longer have sufficient kinetic energy to overcome interparticle attractive forces.

• The freezing point of a solution is always lower than that of the pure solvent.

Freezing Point Depression (cont.)

• Solute particles interfere with the attractive forces among solvent particles.

• A solution's freezing point depression is the difference in temperature between its freezing point and the freezing point of the pure solvent.

ΔTf = iKfm

• where ΔTf is the freezing point depression, i is the van’t Hoff factor, Kf is the freezing point depression constant, and m is molality.

Freezing Point Depression (cont.)

Osmotic Pressure

• Osmosis is the diffusion of a solvent through a semipermeable membrane.

Osmotic Pressure (cont.)

• Osmotic pressure is the amount of additional pressure caused by water molecules that moved that moved into the concentrated solution.

• Osmotic pressure of a solution is always greater than its pure solvent.

Section 14.1 Types of Mixtures

Key Concepts

• The individual substances in a heterogeneous mixture remain distinct.

• Two types of heterogeneous mixtures are suspensions and colloids.

• Brownian motion is the erratic movement of colloid particles.

• Colloids exhibit the Tyndall effect.

• A solution can exist as a gas, a liquid, or a solid, depending on the solvent.

• Solutes in a solution can be gases, liquids, or solids.

Section 14.2 Solution Concentration

Key Concepts

• Concentrations can be measured qualitatively and quantitatively.

• Molarity is the number of moles of solute dissolved per liter of solution.

• Molality is the ratio of the number of moles of solute dissolved in 1 kg of solvent.

Section 14.2 Solution Concentration (cont.)

Key Concepts

• The number of moles of solute does not change during a dilution.

M1V1 = M2V2

Section 14.3 Factors Affecting Solvation

Key Concepts

• The process of solvation involves solute particles surrounded by solvent particles.

• Solutions can be unsaturated, saturated, or supersaturated.

• Henry’s law states that at a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid.

Section 14.4 Colligative Properties of Solutions

Key Concepts

• Nonvolatile solutes lower the vapor pressure of a solution.

• Boiling point elevation is directly related to the solution’s molality.

∆Tb = Kbm

• A solution’s freezing point depression is always lower than that of the pure solvent.

∆Tf = Kfm

• Osmotic pressure depends on the number of solute particles in a given volume.