lesson 3: solubility equilibria

Chemical equilibrium

State in which the rates of the forward and reverse reactions are equal. Concentrations remain constant.

Equilibrium

Is a dynamic process and there is no net change in the number of molecules of reactants or products.

Equilibrium expression

Products raised to coefficients (if any) over reactants raised to coefficients (if any)

aA + bB ⇆ cC + dD; where small letters are the coefficient

Equilibrium constant allows us to predict:

• The extent of reaction when equilibrium is established.

• The direction of net reaction from a given set of initial concentrations.

• Whether a given set of concentrations represents an equilibrium condition.

K > 1 where K is equilibrium constant

Reactant favored

K = 1 where K is equilibrium constant

Value of K at equilibrium

K < 1 where K is equilibrium constant

Product favored

Equilibrium constant for new reaction

Inverse of the original reaction

Homogenous equilibrium

All products and reactants are in the same physical state.

Heterogenous equilibrium

Reactants and products in more than one phase. Does not depend on the amounts of pure solids or liquids.

Le Châtelier's Principle

States that if an external stress (temperature, pressure, concentration, volume) is added to a system, the equilibrium shifts to a new position.

Concentration - equilibrium

• Addition of reactant shifts the equilibrium to the right

• Addition of products shifts the equilibrium to the left

• USE SEASAW METHOD

Pressure - equlibrium

• Increase shifts the equilibrium to the fewer gas molecules.

• Decrease shifts the equilibrium to the side with more gas molecules

Increase in pressure - pressure equilibrium

• Equilibrium shifts in reverse

• Favors the side of fewer moles of gas

Increase in temperature - equilibrium

• Exothermic reaction (heat released, system absorbs heat)

• Favors reverse reactions

• ONLY FACTOR THAT HAS EFFECT ON K

Decrease in temperature - equilibrium

• Endothermic reaction (heat absorbed, system releases heat)

• Favors forward reactions

• ONLY FACTOR THAT HAS EFFECT ON K

Catalysts

Speed up or retard the rate of equilibrium.

Ksp

Equilibrium constant for a sparingly soluble salt

• Lower = more soluble

Reaction quotient

Also called the ion product. Gives the same ratio the equilibrium expression gives, but for a system that is not at equilibrium.

Q = K

System is at equilibrium state

Q > K

Too much product (precipitate will form), equilibrium shits to the reactant side

Q < K

Too much reactant (unsaturated solution), equilibrium shifts to the product side.

Common ion effect (increase in common ion decreases in solubility)

Solubility of sparingly soluble species is reduced. Inverse relationship with solubility. A salt is less soluble if one of its __ is already present in the solution.

Effect of pH

Sparingly soluble from weak acids tend to be more soluble in an acidic solution. Inverse relationship with solubility. More acidic = more soluble

Effect of complexation (ion pairs)

Formation of complex ions increases the solubility of a sparingly soluble salt. Directly proportional relationship with solubility.

Salt effect

Diverse ion increase the Ksp value. In the presence of an inert electrolyte, activities of ions decrease. Adding an inert electrolyte increases the solubility of a sparingly soluble salt. Directly proportional relationship with solubility.

Electrolytes

Create an ionic atmosphere that decreases the attraction between ions in the solution.

Effect of electrolytes

To reduce the tendency for salts to come together thereby increasing solubility.

Ionic strength

Measure of total concentrations of ions in a solution.

Gilbert Newton Lewis

Introduced the concept of activity and formed a model in acids and bases where electrons are involved.

Activity coefficient

Measures the deviation of behavior from ideality. If it’s 1, behavior is ideal and first equation would be correct.

Peter Debye and Erich Huckel

Quantified that an expression permits the calculation of activity.

Interpolation

Estimation of a number that lies between two values in a table.

Higher charge

The lower the activity coefficient (__ and coefficient have inversely proportional relationship).

Higher ionic strength

The lower the activity coefficient (__ and coefficient have inversely proportional relationship).

Salting out effect

Increase of activity coefficients with ionic strength (𝜇≥ 1M)

Svante August Arrhenius

Swedish chemist proposed two classification of compounds; acids and bases.

Arrhenius acid

Increases the concentration of hydronium ions (H+)

Arrhenius base

Increases the concentration of hydroxide ions (OH-)

Thomas Martin Lowry and Johannes Nicolaus Brønsted (Brønsted-Lowry theory)

Proton theory of acids and bases. Introduced independently in 1923.

Brønsted-Lowry acid

Proton donor

Brønsted-Lowry base

Proton acceptor

Conjugate acid

Species that forms after a base accepts a proton

Conjugate base

Species that forms after an acid donates a proton

Lewis acid

Electron pair acceptor

Lewis base

Electron pair donor

Adduct

Compound with coordinate covalent bond, both electrons are provided by only one of the atoms.

Ion product constant of H2O at 25°C

Kw = 1.0×10^-14 M

Calculating for pH

• = -log[H+]

• = -log[A_H+]

• = -log[H+]γ_H+

• = 14 - pOH or 14 - (-log[OH-])

Strong acids and bases

Species that completely ionize in water

Weak acid and weak base or weak electrolyte

Partially ionized (<5%) in aqueous solution.

Stronger acid

Product is weaker conjugate base

Weaker acid

Product is stronger conjugate base

Equilibrium constant expression for acid

Lower equilibrium constant for this, the weaker it is.

Percent ionization

If less than 5%, ionization is negligible.

Stronger base

Product is weaker conjugate base

Weaker base

Product is stronger conjugate acid

Equilibrium constant expression for base

Lower equilibrium constant for this, the weaker it is.

Salt that produce neutral solutions

From strong bases and strong acids, they do not hydrolyze and remain neutral at 7.

Hydrolysis of salt in basic solutions

Cation of a strong base + anion of a weak acid. Anion will accept proton from the water.

Hydrolysis of salt in acidic solutions

Cation of a weak base + anion of a strong acid. Cation is a weak acid that will form H3O+

Hydrolysis of salt

If both anion and cation can react with water, the solution will depend on the greater equilibrium constant value. (i.e. ka > kb, meaning solution is acidic)

Monoprotic acid

Contains one proton

Diprotic acid

Contains two protons. The first proton is the easiest proton removed compared to the successive proton.

Triprotic acid

Contains three protons. The first proton is the easiest proton removed compared to the successive proton.

Polyprotic acid

Contains more than one proton. The first proton is the easiest proton removed compared to the successive proton.

Amphoteric or amphiprotic

A substance that can react as both acid or base.

Buffer solutions

• Resist change in pH on a small dilution

• Contains a weak acid or weak base and their salts

Buffering capacity (BC)

Measure of the resistance of a buffer to pH.

Maximum buffering capacity

When concentration of base is equal to concentration of acid.

pH calculation of buffer, Henderson Hasselbalch equation

In choosing a buffer, seek one whose pKa is as close as possible to the desired pH.

Charge balance

Sum of positive charges equals the sum of the negative charges in a solution.

Mass balance

Quantity of all species in a solution containing a particular atom must equal to the amount of that atom delivered in a solution.

Systematic treatment of equilibrium

Write as many independent equations as there are chemical species in the system. Includes:

• equilibrium constant expressions

• Mass balance equations

• Single charge balance equation