CHEM 120 Chapter 13

Mixtures

2+ components

-Homogenous/heterogenous

Homogenous mixtures

• Solutions

  • Have a uniform appearance

  • Can have all 3 states as a base

  • Can have multiple states in one solution

  • Water base = “aqueous”

A clear solution is ___ not ____

see through, not colorless

heterogenous mixture

can tell the different components apart

ICLICKER

Which of the following is NOT true for solutions?

A) Solutions have a uniform appearance

B) Solutions are a mixture of two or more substances

C) The substance in largest mole amount is the solvent.

D) A mixture is only a solution if the solvent is water.

E) The substance(s) in lesser mole amount is/are the solute(s).

D.

Solvent

does the dissolving

solute

is the dissolved

Solubility

amount of solute that can dissolve per amount of solvent at a given

temperature.

• Units: g/L, M, % mass, mass/volume etc.

• Like dissolves like

• Solute/solvent particle interactions create INTERMolecular forces

Intermolecular forces

• Ion-dipole (ion-polar)

• Dipole-dipole

  • Hydrogen bond (H - N,O,F)

• Ion-induced dipole (ion-nonpolar)

• Dipole-induced dipole (polar-nonpolar)

• Dispersion (nonpolar-nonpolar)

Water and solution formation/Intermolecular forces

Water is a dipole.

The dipole-dipole attraction/ H bonds between water molecules

Ion-dipole between ions and water

• Creates Hydration shells/ spheres (water molecules that surround individual ions)

ion-ion between ions

Lattice enthalpy

ion-ion attraction energy

-Decreases with larger atom size

-increases with larger charge

Enthalpy of hydration (water)/solvation enthalpy (not water)

ion-dipole attraction energy

ICLICKER

B.

“organic” = fats, etc. nonpolar

• nonpolar tail on Vitamin A

• multiple polar groups (OH) and dipole (O=C) on vitamin C)

Dual polarity

molecules can have polar and nonpolar ends

form micelles (Polar ends out, nonpolar ends in) which can take in nonpolar molecules into liquid state

miscibility

ability for a solute to dissolve in all proportions with the solvent

solubility in alcohols in water and in hexane

ICLICKER

Which of the following paired liquids are immiscible in each other?

A) CH3CH2OH and H2O (ethanol and water)

B) NH3 and CH3CH2OH (ammonia and ethanol)

C) CH3CH2CH2CH2CH2CH3 and CH3CH2OH. (hexane and ethanol)

D) CH3CH2CH2CH2CH2CH3 andCH3CH2CH2CH2CH2CH2CH2CH3 (hexane and octane)

E) CH3CH2(OH)CH3 and H2O (isopropanol and water)

None, all are miscible

enthalpy

heat flow at constant pressure (state function- start and end)

denoted by delta H

only denoted by q if not at constant pressure

thermodynamics of solution formation

Sum of enthalpy changes

Delta H is positive (colder as it needs more energy to absorb) (total energy of solute and solvent is positive)

If dissolving increases, delta H is more negative due to energy release (delta H mix is always less than 0)

How do you know a solution formation is exothermic or endothermic

H final < H initial - exothermic and Hsolution < 0

H final > initial- endothermic and Hsolution> 0

ICLICKER

D.

Delta H (lattice energy) is >/< 0, delta H (hydration of ions) >/< 0

Lattice energy enthalpy is always positive, hydration of ions are always negative

Ion radius relationship with lattice energy and hydration of ions

The larger the size, heats of hydration and lattice energies increase

Entropy

S

amount of energy dispersed at constant temp.

aka disorder or randomness of a system

Absolute entropies are always positive

0 entropy is at 0 Kelvin

order of phases from least to most entropy

1. solid

2. liquid

3. gas

Dissolving and entropy

dissolving FAVORS entropy as it increases it for example, KCl(s) → KCl(aq) is a very cold liquid and increases the number of particles

equilibrium

a steady state

unsaturated solution

solute is dissolved but not at max amount dissolved. there is no solute leftover

saturated solution

max amount of solute dissolved, minimal solute leftover, solution at equilibrium

unit: gSolute/100gSolvent

supersaturated solution

equilibrium releases excess dissolved solute (ex: opening a can of soda releases excess dissolved CO2)

How does temperature affect solid/liquid/gas solutes

As temperature increases, solubility of a solid and liquid solute increases but of a gas solute decreases

As pressure increases, the solubility of a gas in water _

increases (more collisions)

• Henry’s Law

Henry’s Law

Sg= kH x Pgas

kH - Henry’s constant (specific to a gas) (the higher kH is, the higher the solubility)

Pgas - Partial pressure of gas

ICLICKER

B. Remember: as temperature increases, the solubility of a gas solute decreases

PRACTICE

Before: 0.17 M CO2

After: 1.36 × 10^-5 M

Molarity

M

• concentration at a specific temperature

• (solute) mol/ (SOLUTION volume) L

molality

m OR m

• concentration independent of temp

• (solute) mol / (SOLVENT mass) kg

parts by mass

mass of solute/mass of solutionp

parts by volume

volume of solute/volume of solution

mole fraction

X (percentage before x 100)

(solute) mol / ((solute) mol + (solvent) mol)

PRACTICE

molality- 8.388 m

mole fraction- 0.1313

% by mass- 32.89%

PRACTICE

1) 0.5 m

2) 0.01 m

PRACTICE

At 60 degrees: NH4Cl and NaCl are saturated and RbCl is unsaturated

At room temp: all are saturated

ICLICKER

A. 35 is less than 50

Colligative properties

changes in solution properties that depend on the number of solute particles, not the identity of the solute

• boiling point elevation, freezing point depression, osmotic pressure, vapor pressure lowering

• Dissolving increases boiling point of the solution

electrolytes

solutes that carry a charge through a solution

strong electrolytes

strong acids/bases, soluble ionic compounds

are completely ionized and dissociated

weak electrolytes

weak acids/bases, partially ionized and dissociated

non-electrolytes

sugar, ethylene glycol, ethanol, etc. don’t dissociate or ionize in water

dissociate

forming cations/anions

• strong: 2

• weak: 1-2

• non: 1

Raoult’s Law

• property of non-electrolytes

Vapor pressure lowering

VP of a solvent is proportional to mol fraction (as % of the solvent increases, then the final VP is proportional to it

volatile

ability to evaporate

why do nonvolatile solutes lower VP

decreases surface area for evaporation

PRACTICE

23.3 torr

boiling point elevation

Vapor pressure = environmental pressure

solutions always boil at a temp higher than the pure solvent (as molality increases, so does the change in vapor pressure)

1 atm pressure = 100 degrees Celsius (normal boil point)

to compensate, the solution must be heated to a higher temp

freezing point depression

solutions always freeze at a temp lower than the pure solvent

• formation of a lattice by pushing out dissolved solutes

solute disrupts freezing process, solution must be cooled to a lower temp to freeze solvent

PRACTICE

-2.3 degrees Celsius

osmotic pressure

pressure to keep solute from going through a semipermeable membrane