Oceans A - Solutions, Enthalpy Changes, Acid-base Equilibria

Define ionic bonding

The electrostatic attraction between positive and negative ions

What two processes occur when an ionic solid dissolves

• The ionic lattice breaks into gaseous ions

- Energy is required to overcome forces of attraction between ions

- This is the reverse of lattice enthalpy

• Gaseous ions dissolve in water (hydration) due to the attraction between ions and polar water molecules (ion-dipole bonds)

What factor impacts if an ionic solid will dissolve

Solid will dissolve if more (or similar ) energy is released when bonds form, than is needed to break bonds

• Bonds broken = ionic bonds in solid and hydrogen bonds in water

• Bonds formed = ion-dipole bonds between ions and polar water molecules

What is Lattice Enthalpy and what does its magnitude depend on

The enthalpy change on formation of 1 mole of an ionic solid from gaseous ions

Is always exothermic (negative) because ionic bonds are being made

Magnitude (more exothermic) depend on strength of ionic bonds which depends on the charge and size of ions involved

What is Enthalpy change of hydration and what does its magnitude depend on

The enthalpy change on dissolving 1 moles of gaseous ions in water

Always exothermic (Negative) as ion-dipole bonds are being made between ions and polar water molecules

Ions become hydrated

Magnitude (more exothermic) depends on ability of the ions to attract the polar water molecules forming ion-dipole bonds

The attraction is strongest with ions with high charge density

What is enthalpy change of solution

The enthalpy change on dissolving 1 moles of a solute forming an infinitely dilute solution

Can be exothermic or endothermic

Lattice enthalpy, enthalpy change of hydration and enthalpy change of solution can be linked with a Born Habor cycle. Draw a Born Habor Cycle for calcium chloride (CaCl2)

If enthalpy change of a solution is exothermic, what will the solubility of the solid be

• Almost always soluble

• Providing enthalpy change of hydration is more exothermic than lattice enthalpy, the enthalpy change of solution will be negative and the solid will dissolve

• The energy needed to break up the ionic lattice is more than compensated for by the exothermic enthalpy of hydration

If enthalpy change of a solution is endothermic, what will the solubility of the solid be

• Solubility depends on how endothermic the enthalpy of solution is 

• Insoluble - enthalpy of solution is large and endothermic (lattice enthalpy must more exo than enthalpy of hydration)

• Usually soluble - enthalpy of solution is small and endothermic. Heat energy taken in from surroundings 

Use enthalpy changes to explain why Group 2 hydroxides increase in solubility down the group

• Controlled by lattice enthalpy of M(OH)2

• Lattice enthalpy decreases (less exo) down the group because charge density decreases

• Enthalpy of hydration also decreases (less exo) down the group due to lower charge density

• Lattice enthalpy decreases more rapidly, so as cation size increases, Enthalpy of hydration becomes > lattice enthalpy and solubility increases

According to Bronsted-Lowry Theory, an acid is _____ and a base is ____

Acid - proton donor

Base - proton acceptor

What is the difference between an alkali and a base

A base - a proton acceptor

An alkali - Dissolves in water to form OH- ions

All alkalis are bases but not all bases are alkaline

Acids contain _____ atoms bonded ______ to an _________ atom. They are polar molecules

In water, acid molecules collide with ______ water molecules and a _____ bond forms between the O on water and the H of the acid.

The _____ ion that is formed is called a __________ ion. These are present in every dilute acid and responsible for the behaviour of all acids.

In terms of equations, we simplify and use ____ instead

Acids contain hydrogen atoms bonded covalently to an electronegative atom. They are polar molecules

In water, acid molecules collide with polar water molecules and a dative bond forms between the O on water and the H of the acid.

The H3O+ ion that is formed is called a hydroxonium ion. These are present in every dilute acid and responsible for the behaviour of all acids.

In terms of equations, we simplify and use H+ instead

What is a strong acid/base

• A strong acid fully dissociates into ions in solutions

E.g. HCl → H+ + Cl-

• A strong base fully dissociates into ions in solutions

E.g. NaOH → Na+ + OH-

What is a weak acid/base

• A weak acid partially dissociates into its ions in solutions

E.g. CH3COOH <=> H+ + CH3COO-

• A weak base partially dissociates into its ions in solutions

E.g. NH3 + H2O <=> NH4+ + OH-

Acid + Metal →

Salt + H2

Acid + Base →

Salt + Water

Acid + Metal Carbonate →

CO2 + Salt + H2O

Acid + Alkali →

Salt + Water

What is the general equation for an acid + base equilibrium

HA + H2O <=> H3O+ + A-

How do acid base indicators work

Indicators are weak acids

HA <=> H+ + A-

Either HA or A- must be coloured

Addition of acid causes concentration of H+ to increase

Equilibrium shifts to LHS

The colour of HA appears

Addition of alkali cause H+ to react and therefore to decrease in concentration

Equilibrium shifts to RHS

The colour of A- appears

What is a conjugate acid-base pair

a conjugate acid–base pair consists of two substances that differ only by the presence of a proton (H⁺). A conjugate acid is formed when a proton is added to a base, and a conjugate base is formed when a proton is removed from an acid.

Will conjugate bases be strong or weak if the acid is strong

• Conjugative base will be weak

• Acid is strong so dissociates fully

• Reactions complete

• Equilibrium fully to RHS, reverse reaction will not occur

• Conjugative base does not act as a base in its reaction (reverse reaction)

Will conjugate bases be strong or weak if the acid is weak

• Conjugate base will be strong

• Weak acid partially dissociates

• Reaction doesn’t go to completion

• Equilibrium position to left as reverse reaction occurs

• Conjugate base acts as a base in its reactions (reverse)

What is the expression for the Acid Dissociation Constant, Ka

Ka = [H+] [A-] / [HA]

What is the only factor that can change Ka

Temperature

For strong acids:

• Equilibrium is to the ______

• Because the dissociate _____ in water

• Therefore the magnitude of Ka would be _____

• And the magnitude of pKa would be _____

For strong acids:

• Equilibrium is to the RHS

• Because the dissociate fully in water

• Therefore the magnitude of Ka would be Large

• And the magnitude of pKa would be Small

For weak acids:

• Equilibrium is to the _____

• Because the dissociate _____ in water

• Therefore the magnitude of Ka would be _____

• And the magnitude of pKa would be _____

For weak acids:

• Equilibrium is to the LHS

• Because the dissociate partially in water

• Therefore the magnitude of Ka would be Small

• And the magnitude of pKa would be Large

What is the equation that links Ka with pKa

What is the equation for the ionic product of water, Kw

What is the value for Kw at room temp (on data sheet)

Kw = [H+] [OH-]

At room temp Kw = 1.00 × 10^-14 mol2 dm^-6

Why can pure water carry electrical current

Pure water will conduct electric current due to the small amount of dissociation

H2O <=> H+ + OH-

How can pH be calculated using the H+ concentration

Use water as an example which has a H+ concentration of 1 × 10^-7 mol dm-3

PH = -log10 [H+]

Water pH = -log10 (1 × 10^-7)

= 7

How do acidic buffers work, use ethanol acid and sodium ethanoate as an example

A buffer doesn’t stop the pH from changing completely, but it minimises it

Acidic buffers are made by mixing a weak acid with its conjugate base (salt)

In this example, sodium ethanoate fully dissociates in water

Ethanoic acid only partially dissociates in water

Therefore the buffer solution contains lots of CH3COO- and lots of CH3COOH

On addition of acid

• The equilibrium position shifts to the LHS

• CH3COO-(aq) + H+(aq) → CH3COOH

• CH3COO- reacts with additional H+ so the change in pH is minimised

On addition of alkali

• H+ ions are removed when they react with OH- to form water

• The equilibrium position shifted to the RHS to replace H+

• OH- + H+ → H2O

• Change in pH is minimised

Overall

PH change is counteracted as the concentrations of weak acid and conjugate base are both large

How do buffers play a role in maintains the blood plasma pH

H2CO3(aq) <=> H+(aq) + HCO3(aq)

NaHCO3 (aq) → Na+(aq) + HCO3-(aq)

On addition of an acid

• Concentration of H+ increases

• H+ reacts with HCO3-

• Equilibrium over to the LHS

On addition of an alkali

• Concentration of OH- increases

• OH- reacts with H+ to form H2O

• Equilibrium over to the RHS

What happens to the pH of a buffer if you dilute it

• Remains unchanged

• If the buffer solution is diluted, the concentration of both acid (HA) and conjugate base (A-) will change by the same factor

• On dilution, the value of [A-] / [HA] will be unchanged

• As Ka is a constant, [H+] must be unchanged too

What happens to the pH of an acid when you dilute it

• The concentration of H+ decreases

• As the [H+] concentration of water is 1 × 10^-7 mol/dm³, the more water that is added to the acid, the closer [H+] becomes 1 × 10^-7

• The pH of the acid tends towards 7