Chem Notes Ch 5

Question 1

Solute

Answer 1

Substance present in a smaller proportion

Question 2

Solvent

Answer 2

Substance present in a larger proportion

Question 3

Aqueous solution

Answer 3

Solvent is water

Question 4

Solubility

Answer 4

Amount of solute that will saturate a particular solvent; rate of dissolution is equal to the rate of precipitation; The solid form and the dissolved form are in dynamic equilibrium.

For solids in liquids, solubility increases with temperature. Gases are more volatile at higher temperatures, which causes them to be less soluble in liquids. A gas is more soluble in a liquid when the pressure is increased.

Question 5

Phase solubility rules - solids in liquids

Answer 5

Increased temperature = increased solubility

Question 6

Phase solubility rules - gases in liquids

Answer 6

Increased temperature = decreased solubility

Increased pressure = increased solubility

Question 7

Crystallization

Answer 7

Process of crystal formation when a dissolved solute comes out of a solution

Question 8

Supersaturated

Answer 8

Solutions contain more solute than found in a saturated solution. They are formed when manipulating temperature or pressure.

Question 9

Salt solubility rules

Answer 9

1. All salts of Group 1 metal cations (Li(+), Na(+), K(+), Rb(+), Cs(+)) and ammonium (NH4(+)) are
   soluble.
2. All salts of nitrate (NO3(-)), perchlorate (CIO4(-)) and acetate (C2H3O2(-)) are soluble.
3. All salts of silver (Ag(+)), lead (Pb(2+)), and mercury (Hg(2+)) are insoluble.
4. All salts of hydroxide (OH(-)), carbonate (CO3(2-)), phosphate (PO4(3-)), and sulfide (S(2-)) are
   insoluble.
5. When a salt consists of a “soluble ion” and an “insoluble ion”, it is SOLUBLE in water
   (ex. NaOH = soluble).

Question 10

Molarity (M)

Answer 10

Concentration of a solution expressed in moles of solute per volume (in liters) of solution. Dependent of both temperature and pressure.

Question 11

Molality (m)

Answer 11

Concentration in terms of moles of solute per mass (in kilograms) of solvent. Independent of both temperature and pressure.

Question 12

Mole fraction

Answer 12

Moles of one component of a solution divided by the total number of moles in solution. It is unit-less because it is a ratio

Question 13

Normality (N)

Answer 13

The number of mole equivalents (n) per litre of solution. Used for acids and bases.

Formula: (Number of moles)*(Molarity)

The mole equivalents of an acid or base are calculated by determining the number of H+ or OH- ions per molecule. For instance, the mole equivalents of H2SO4 is two because dissolution of sulfuric acid releases two H+ ions into
solution.

Question 14

Colligative Properties

Answer 14

Dependent on the number of solute particles in the solution.

1. Vapor-pressure depression
2. Boiling-point elevation
3. Freezing-point depression
4. Osmotic pressure

Question 15

Vapor-Pressure Depression

Answer 15

A colligative property. The pressure exerted by the gaseous phase of a liquid that evaporated from the exposed surface of the liquid.

A higher vapor pressure corresponds to a lower boiling point and the more easily a substance will evaporate.

If a solution contains a dissolved solute, the solute molecules are in a state of high entropy, decreasing the vapour pressure, making it more difficult for the solution to boil and requires more energy.

Question 16

Boiling point

Answer 16

Temperature at which the vapor pressure of the solution is equal to the atmospheric pressure.

At sea level, atmospheric pressure is 760 mmHg

Question 17

Raoult’s Law

Answer 17

The vapor pressure (PA) of a solution with a non-volatile solute is equal to the pure vapor pressure of the solvent (P°A) multiplied by the mole fraction of the solvent (XA).

The mole fraction of the solvent is the moles of solvent divided by the total number of moles in solution.

Since the presence of solute causes XA to be smaller than 1, PA is lower than P°A

Question 18

Boiling-Point Elevation

Answer 18

A colligative property. The increase in boiling point of solution.

Δ(Tb) = (kb)(i)(m)

kb is the solvent’s boiling-point elevation constant, i is the solute’s van ’t Hoff factor or the number of particles the solute breaks into when it dissolves, and m is the molal concentration of the solution.

Question 19

Freezing-Point Depression

Answer 19

A colligative property. Liquid will be less able to achieve a solid state when a solute is present and the freezing point will decrease.

Δ(Tf) = -(kf)*(i)(m)

kf is the solvent’s freezing-point depression constant, i is the solute’s van ’t Hoff factor, and m is the molal concentration of the solution

Question 20

Osmotic Pressure

Answer 20

A colligative property. The pressure it would take to stop osmosis (net movement of water across a semipermeable membrane from an area of low solute concentration to a region of higher solute concentration) from happening.

van’t Hoff equation can be used to compute osmotic pressure: Π= (M)(i)(R)*(T)

Π is osmotic pressure in atm, M is molarity of the solution, i is the van’t Hoff factor, R is
the universal gas constant of 0.082, and T is the temperature in Kelvin.

Question 21

Non-Colligative Properties

Answer 21

Depend on the identity of the dissolved molecules and the solvent. Ex. colour, viscosity, surface tension, and solubility.

Question 22

Dissolution

Answer 22

When the solute breaks up from a larger crystal of molecules into much smaller groups or individual molecules. Caused by interactions between the solute and solvent. Intermolecular forces that contribute to dissolution include dipole-dipole, Van der Waals, ion-dipole, and hydrogen bonding.

Question 23

From weakest to strongest, the list of intermolecular forces

Answer 23

London dispersion, to dipole-dipole, to hydrogen bond, to ion-dipole.

Question 24

Polarity

Answer 24

Solutes dissolve best in solvents with similar intermolecular forces. Solutes whose polarity matches that of the solvent will most likely be soluble.

The intermolecular forces being broken in the solute are being replaced by equal (or stronger) intermolecular forces between the solvent and the solute.

Question 25

Electrolytes

Answer 25

Used to make conductive solutions. Pure water is an example of a solution that does not conduct an electrical current.

A strong electrolyte dissociates completely into its constituent ions. (Ex. NaCl and KI, which are both ionic compounds. Molecular compounds with highly polar covalent bonds that dissociate into ions when dissolved, such as HCl in water.)

A weak electrolyte would ionize or hydrolyze incompletely in aqueous solution, with only a portion of the solute present in ionic form. (Ex. Acetic acid, ammonia, and HgCl2. Nonpolar gases and organic compounds, such as sugar, do not ionize at all in aqueous solution.)

Question 26

Dielectric constant

Answer 26

Determines polarity of solvent, higher value is more polar