Oceans

lattice enthalpy

energy change when one mole of solid is formed by the coming together of separate gaseous ions

All lattice enthalpies are

negative

What amount of energy do you need to put in to break down a lattice?

-∆_LE H

Lattice enthalpies become more negative

• ionic charge increases

• ionic radii decreases

Lattice enthalpy is more negative

Why do ions with higher charge densities have more negative lattice enthalpies?

• higher charges attract more strongly

• closer ions attract more strongly

• ions with smaller radii can come closer together and attract more strongly

Stronger attraction

more negative lattice enthalpy

Why are water molecules polar and behave as a dipole?

• covalent bonds are polar

• bent shape means molecule is polar

Hydration

polar water molecules move to surround ions

negative end of dipole faces positive ions

positive end of dipole faces negative ions

If water dipoles bond weakly, the ion is

not extensively hydrated

If water dipoles bond strongly, the ion is

extensively hydrated

Higher charge density means for a hydrated ion

• more water molecules attracted

• bigger hydrated ion

Why may hydration overcome lattice enthalpies?

• Strong ion–dipole attractions

• Small, highly charged ions
  Ions with high charge density attract water molecules very strongly → more -hydration enthalpy

• energy released during hydration is greater than the energy needed to break the lattice
  → overall energy change is negative
  → dissolution is energetically favorable

Breaking hydrogen bonds in water requires

energy

Enthalpy change hydration

enthalpy change for formation of a solution of ions from one mole of gaseous ions

Enthalpy change of hydration depends on

concentration

Why do we use very dilute solutions for enthalpy change of hydration?

assume interactions between ions are negligible

Most exothermic enthalpy change of hydration

• greatest charge

• smallest radii

Enthalpy change of solution

enthalpy change when one mole of solute dissolves to form a very dilute solution

Equation of enthalpy change solution

∆_solutionH= ∆_hydH(cation) + ∆_hyd(anion) - ∆_LEH

When does a solution normally dissolve?

∆_solH is negative

when does a solute not dissolve?

∆_solH is large and positive

Why does an ionic substance dissolve when reaction appears energetically unfavourable?

∆_solH is slightly positive

Enthalpy diagram for an ionic substance in a non-polar solvent

What happens to energy absorbed by greenhouse gas molecules?

1. Increase vibrational energy of molecules so bonds vibrate more and transfer kinetic energy to other molecules by collision, increasing temperature.

2. Some radiation is re-emitted to Earth and space

IR window

wavelengths of IR water vapour does not absorb so escape into space

Carbon dioxide _____________ wavelengths in the IR window

absorbs

The higher the temperature, the _________ the level of water vapour.

higher

Bronsted Lowry theory of acids and bases

Acids are H⁺/ proton donors

Bases are H⁺/ proton acceptors

conjugate base

remains after acid donates H⁺

conjugate acid

forms when base accepts H⁺

Conjugate acid base of HCl (aq)

Cl⁻

Conjugate acid of NH₃

NH₄⁺

Water acting as a base with HCl

HCl (aq) + H₂O (l) → H₃O⁺ (aq) + Cl⁻ (aq)

Water acting as an acid with NH₃

H₂O + NH₃ → OH⁻ + NH₄⁺

pH

-log [H⁺]

low pH

high [H⁺]

Strong acids

strong tendency to donate H⁺

dissociate completely in aqueous solution

Reaction strong acid HA with water

HA (aq) + H₂O (l) → H₃O⁺ (aq) + A⁻ (aq)

Examples strong acids

HCl

H₂SO₄

weak acid

weak tendency to donate H⁺

does not dissociate completely in water

Dissociation weak acid HA

HA (aq) ⇌ H⁺ (aq) + A⁻ (aq)

What does the equilibrium sign show?

a significant amount of HA molecule is present

The further the equilibrium lies to the right-hand side, the ________ the acid

stronger

How to calculate pH of strong acid?

-log [HA= H⁺]

Assumptions for calculating pH of weak acids?

1. [H⁺] = [A⁻] (including H⁺ from water since very small amount)

2. [HA] in equilibrium = [HA] put into solution (very small fraction H⁺)

How to calculate pH of weak acid?

1. Acidity constant

2. -log[H⁺]

Acidity constant

How to calculate pKa?

pKa= -log Ka

How to calculate pH of strong alkali?

1. Ionisation product of water

2. -log[H⁺]

Equation Kw water?

Kw = [H⁺] [OH⁻] / [alkali]

[alkali] = [OH⁻]

alkali in excess so ignore

Kw= [H⁺][OH⁻]

The larger the Ka value

the stronger the weak acid

The larger the pKa value

the weaker the weak acid

Buffer

resists changes in pH on addition small amounts of acids and alkalis

What are buffer solutions made from?

1. Weak acid + salt

2. Weak base + salt

Assumption about buffers made from weak acid and salt?

1. All A⁻ come from salt

2. Almost all HA put into buffer remain unchanged

What happens in a buffer when H⁺ are added?

A⁻ from salt react with H⁺ to form HA and water

removes excess H⁺ and pH re-established

What happens when OH^- are added in a buffer?

H⁺ ions are removed to form water

H⁺ ions are regenerated by HA and pH is re-established

How does a solution of H+ and A- ions act as a buffer?

HA → H+ + A-

1. When H+ is added, equilibrium shifts to the left to form acid and remove excess H+

2. When OH- is added, equilibrium shifts to the right to replace H+

3. Both H+ and OH- are in excess

Ka buffer using assumptions

What does the [H⁺] in a buffer depend on?

1. value Ka

2. salt: acid

Why is pH of buffer not affected by dilution?

concentration of salt and acid are reduced equally

K_sp

Solubility product

What does K_sp represent?

conditions for equilibrium between sparingly soluble solid and saturated solution

How to calculate K_sp for reaction AB → A + B?

[A]^a[B]^b

Ion product, Q

state of the solution right now

How to find Q?

[conc at that moment]

Why does precipitate form when Q>Ksp?

ions present in higher concentration than the solution can hold,

extra solid precipitates out.

If value Q in excess of Ksp?

precipitate out

If Q smaller than Ksp?

no precipitate

If Q=Ksp?

saturated

Solubility

maximum amount of a substance that dissolves in water to form a saturated solution

mol/L

How to calculate solubility of HA → H+ + A-

[H+]=[A-]

Ksp= s x s = s²

s= Ksp^1/2

Entropy

measure of number of ways of arranging molecules and distributing energy

When does a collection of molecules have higher entropy?

• spread out more

• energy is shared among more molecules

How to calculate total entropy change?

∆_totalS= ∆_sysS + ∆_surrS

How to calculate enthalpy change of surroundings?

-∆H/ T

When is reaction feasible?

positive ∆_totalS

What kind of value is ∆_totalS for spontaneous reactions?

positive

What happens when ∆_totalS is 0?

equilibrium is constant

Kc= 1

When are positive ∆_totalS reactions not spontaneous?

large activation enthalpy