unit 4: solutions and solubility

Solution:

homogeneous mixture

Solute:

substance being dissolved, part of the solution present in the smaller amount

Solvent:

part of the solution present in the greatest amount, does the dissolving

Miscible:

liquids that mix completely

Immiscible:

unable to mix

SOLUBILITY AND FORCES BETWEEN PARTICLES

• In order to dissolve:

  • the solute particles must be attracted to the solvent particles

  • the intermolecular forces between solute and between solvent  particles need to be broken

1. Types of Solute & Solvent Particles

• Ionic and polar solutes will dissolve in polar solvents because particles of both are charged

• Polar solutes will dissolve in polar solvents

• Non-polar solvents dissolve non-polar solutes due to similar intermolecular forces (no full or partial charges) 

• Recall: ∆EN of non-polar molecules is < 0.5 (0.4 and below), polar molecules is between 1.6-0.5, and  ionic compounds is 1.7 and above

This relationship is summarized in the expression: 

“Like dissolves like”

“Like dissolves like”

2. Temperature (for solids)

For Solids:

• Increases in temperature causes increased solubility as a higher in temperature causes:

  • Spaces between particles increase resulting in more space for particles of solute to dissolve

  • Solvent particles have greater kinetic energy which results in more frequent and energetic collisions with the solute

temperature for gases

• Increases in temperature causes decreased solubility 

  • Molecules in gaseous state have higher kinetic energy than those dissolved in the solvent

  • Increasing temperature provides energy for gas molecules to escape solution

3. Pressure

• Pressure is force per unit area

• No effect on solubility of solids or liquids

• Solubility of a gas is directly proportional to the pressure of that gas above the liquid

Increased pressure causes increased solubility of a gas

4. Size

Covalent Compounds

• Increased molecule size (molecular compounds) causes decreased solubility

• Molecules like methanol (CH₃OH) have a non-polar (CH₃) end and a polar (OH) end

• The –OH group predominates and allows the entire molecule to be soluble in water (polar molecule)

• Increasing the size of the non-polar portion decreases solubility

size and ionic compounds

• If the attraction between the ions is very strong, they will be difficult to separate, and won’t be soluble in water

• Solubility usually increases with increased ion size and decreased ion charge (higher, more concentrated charge, stronger attraction between ions, harder to separate, harder to dissolve)

Factors that affect rate of dissolving

How quickly a solute dissolves in a solvent will increase when each of the following is increased:

• Agitation or mixing: increases number of collisions

• Temperature: increased kinetic energy causes more frequent collisions

• Surface area: more solute is in direct contact with the solvent

Solubility

refers to if a substance dissolves in another (the amount of solute that is able to dissolve in a given quantity of solvent)

Saturated

the max amount of solute is dissolved at that temperature

Unsaturated

less than the max amount of solute is dissolved at that temperature

Supersaturated

more than the max amount of solute is dissolved at that temperature

solubility curve

Increasing temperature will increase solubility of solids in a solution

A solubility curve shows the relationship between the solubility of the solute and the temperature of the solution.

• Solubility curves provide information on the solubility of a substance as we change the temperature of solution.

percent concentration

• used for fairly large concentrations

• where the solute is a large proportion of the solution

formula: %conc = amount solute/amount solution x 100

3 ways to get percent concentration

% m/m (percent mass/mass) (ex: g and g)

% m/V (percent mass/volume) (usually g/mL)

% v/v (percent volume/volume) (ex: mL and mL)

ppm and ppb

These units are used to describe the concentration of substances which are present in VERY SMALL AMOUNTS in solution. UNITS MUST MATCH!

molar concentration

• Most useful unit of concentration

• Number of moles of solute per litre of solution

• Aka Molarity (M)

C = Concentration (mol/L or M)

n = moles (mol)

V = volume (L)

C = n/V

dilutions

• When we dilute a solution, we move the solute particles apart  

• No new solute is added, therefore the number of particles stays the same

• The volume changes

• Therefore, the concentration changes

dilution sensory clues

When we dilute a solution, there are sensory clues that tell us the particles have spread out.  These include:

• A lighter colour than the concentrated solution

• A weaker odour if the concentrated solution had an odour

dilution equation

General Equation for DILUTIONS ONLY!!

This equation is NOT for stoichiometry questions!  Here, the substance stays the same!  If the substance changes, you need a mole ratio.

C₁V₁ = C₂V₂

C₁: concentration solution

C₂: diluted solution

properties of ionic compounds

• high melting point

• hard and brittle

• conduct electricity in the dissolved form

• soluble in water, insoluble in cyclohexane

disassociation equations

Dissociation equations show the breaking up of compounds into ions (when they dissolve)

ex: Na₂SO_{4 (S)} → 2Na⁺_(aq) +  SO₄ ^2-_(aq) 

not all ionic compounds are dissolvable

… not ALL ionic compounds completely dissolve

• Some ionic compounds are insoluble (do not dissociate)

The bonds between these ions are strong and are not able to be replaced by water molecules

when are solubility rules used

• Used to predict which chemicals will form precipitates (solid) when reacted.  

• These types of reactions are called precipitation reactions.

write precipitation reactions in 3 steps

1. balanced chemical equation

2. total ionic equation

3. net ionic equation

1. balanced chem equation

ex: Pb(NO₃)_2(aq) + Na₂CO_3(aq) → PbCO_3(s) + 2NaNO_3(aq)

2. total ionic equation

Pb²⁺_(aq) + 2NO₃^-_(aq) + 2Na⁺_(aq) + CO₃^2-_(aq) → PbCO_3(s) + 2Na⁺_(aq) + 2NO_3(aq)

3. net ionic equation

Pb²⁺_(aq) + CO₃^2-_(aq) → PbCO_3(s)

solution stoich steps

1. Balance equation

2. Convert to moles

3. Mole ratio

4. Convert from moles to desired quantity (mass, particles etc)

acids

• Ionize in solution

• Proton donors (Bronsted-Lowry definition)

• Produces hydrogen ions, H⁺_(aq) in water according to Arrhenius definition) 

• The hydrogen ion (H⁺) bonds with a water molecule (H₂O) to make a hydronium ion, H₃O⁺_(aq)

strong acids

• Ionize extremely well

• Completely react with water to make H₃O⁺

strong acids

• Ionize extremely well

• Completely react with water to make H₃O⁺

H₂SO_4(aq) + 2H₂O_(l) → 2H₃O⁺_(aq) + SO₄^2-_(aq)

Weak Acids

• Only a small percentage ionize(most acids are weak acids)

• HC₂H₃O_2(aq) + H₂O_(l) → H₃O⁺_(aq) + C₂H₃O₂^-_(aq)

Bases

• Produces hydroxide ions, OH^-_(aq) (Arrhenius definition)

• Proton acceptors (Bronsted-Lowry definition)

• Bitter and slippery

• Conducts electricity in solution

Strong bases

• Dissociates 100% in aqueous solution

Example

NaOH_(s) → Na⁺_(aq) + OH^-_(aq)

KOH_(s) → K⁺_(aq) + OH^-_(aq)

Weak base

• Ionizes poorly

• Only a small percentage ionize

comparing acids and bases

pH scale

[H₃O⁺] > [OH^-] means concentration of hydronium ions is greater than concentration of hydroxide ions

calculating pH

pH = -log [H₃O⁺]

[ ] means concentration

[H₃O+] means conc. of hydronium or H⁺

calculating pOH

pOH = -log[OH^-]

converting pOH to pH

pH + pOH = 14