Module 5 - Physical Chemistry and Transition Elements

\[5.1.1\]

\[5.1.2\]

[5.1.3] What is a Brønsted–Lowry acid?

A substance that donates a proton.

[5.1.3] What is a Brønsted–Lowry base?

A substance that accepts a proton.

[5.1.3] What are conjugate acid-base pairs?

Conjugate acid-base pairs contain two species that can be interconverted by the transfer of a proton.

e.g. in the dissociation of hydrochloric acid:

HCl ⇌ H⁺ + Cl⁻

In the forward reaction, HCl releases a proton to form its conjugate base, Cl⁻.

In the reverse reaction, Cl⁻ accepts a proton to form its conjugate acid, HCl.

[5.1.3] What is an example of conjugate base pairing?

C₆H₅COOH (pKₐ = 4.19), CH₃CHOHCOOH (pKₐ = 3.86)

C₆H₅COOH (B1) + CH₃CHOHCOOH (A1) ⇌ C₆H₅COOH₂⁺ (A1) + CH₃CHOHCOO⁻ (B2)

the stronger acid donates a proton

[5.1.3] What is a monobasic acid?

An acid with one hydrogen present, e.g. HCl.

[5.1.3] What is a dibasic acid?

An acid with two hydrogens present, e.g. H₂SO₄.

[5.1.3] What is a tribasic acid?

An acid with three hydrogens present, e.g. H₃PO₄.

[5.1.3] Which equation is used to determine the pH of a substance?

pH = -log[H⁺]

[5.1.3] Which equation is used to determine the [H⁺] from its pH?

[H⁺] = 10⁻ᵖᴴ

[5.1.3] What do we assume about strong acids?

We assume that [H⁺] = [HA] as strong acids fully dissociate. The exception is dibasic acids.

[5.1.3] What do we assume about weak acids?

We cannot assume that [H⁺] = [HA] as weak acids only partially dissociate. However, [H⁺] = [A⁻] as each [HA] will dissociate into [H⁺][A⁻].

[5.1.3] What else do we assume about [HA]?

We assume that [HA] at equilibrium is the same as the initial concentration.

[5.1.3] What is the acid dissociation constant?

Kₐ = [H⁺][A⁻] / [HA]

[5.1.3] How can the Kₐ equation be mathematically simplified?

Kₐ = [H⁺]² / [HA]

[5.1.3] How is Kₐ an indicator of acid strength?

Stronger acids have a higher Kₐ value.

[5.1.3] What is the link between Kₐ and pKₐ?

Kₐ and pKₐ both show the strength of an acid.

[5.1.3] What is the equation for pKₐ?

pKₐ = -logKₐ

[5.1.3] How can pKₐ be converted back to Kₐ?

Kₐ = 10⁻ᵖᴷᵃ

[5.1.3] How is pKₐ an indicator of acid strength?

Lower pKₐ = stronger acid.

[5.1.3] What is a strong base?

A substance that accepts protons, and fully dissociates to release all its OH⁻ ions.

[5.1.3] What is the ionic product of water?

The mathematical product of the concentration of hydrogen ions and hydroxide ions.

[5.1.3] What is the Kw equation?

Kw = [H⁺][OH⁻]

[5.1.3] What is the value of Kw at room temperature?

1 × 10⁻¹⁴mol²dm⁻⁶

[5.1.3] Water has a different pH at its boiling point, but remains neutral. Why?

pH decreases as the temperature increases, the equilibrium shifts to favour the forward reaction (endothermic side) and the [H⁺] increases.

Water remains neutral as [H⁺] = [OH⁻].

[5.1.3] What do we assume about strong bases?

We cannot assume [H⁺] = [OH⁻] as we have extra OH⁻ ions; we use the base concentration as the [OH⁻] concentration.

• [H⁺] = Kw / [OH⁻]

[5.1.3] What are conjugate acid-base pairs?

Two species that can be interconverted by the transfer of a proton.

[5.1.3] In the equation: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, what is the conjugate acid and conjugate base?

Base (B1) = NH₃

Conjugate Base (B2) = OH⁻

Acid (A1) = H₂O

Conjugate Acid = NH₄⁺

[5.1.3] In a reaction between two acids, how do we work out the conjugate acid and conjugate base?

The stronger acid is the proton donor, and the weaker acid is the proton acceptor.

[5.1.3] What is a buffer?

Systems which minimise pH changes on addition of small amounts of acid or base.

[5.1.3] What is a weak acid buffer?

A weak acid and a salt of the weak acid, dissolved in water.

[5.1.3] What assumption do we make for buffer calculations?

[H⁺] = Kₐ x [HA]/[A⁻]

[5.1.3] How does equilibrium affect buffers?

HA ⇌ H⁺ + A⁻

• Adding acid will increase [H⁺] so the equilibrium shifts to the left.

• Adding base will shift equilibrium to the right as H⁺ reacts with OH⁻ to make water, causing [H⁺] to decrease.

\[5.2.1\] What is meant by the term **lattice enthalpy,** ΔHₗₑᶿ?

The formation of **1 mole** of **ionic lattice** from its **gaseous ions** under standard conditions. \[***Exothermic]***

\
It is used as a measure of **ionic bond strength**.

\[5.2.1\] What is meant by the term **enthalpy change of atomisation**, ΔHₐₜᶿ?

The energy required for the formation of one mole of gaseous atoms from its element in its standard state. \[***Endothermic***\]

\[5.2.1\] What is meant by the term **first ionisation energy**, ΔHᵢₑ₁ᶿ?

The energy required to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of gaseous +1 ions. \[***Endothermic***\]

\[5.2.1\] What is meant by the term **second ionisation energy**, ΔHᵢₑ₂ᶿ?

The energy required to remove 1 mol of electrons from 1 mol of gaseous +1 ions to form 1 mol of gaseous +2 ions. \[***Endothermic***\]

\[5.2.1\] What is meant by the **first electron affinity**, ΔHₑₐ₁ᶿ?

The energy released when 1 mol of gaseous atoms each acquire an electron to form 1 mol of gaseous 1- ions. \[***Exothermic***\]

\[5.2.1\] What is meant by the **second electron affinity**, ΔHₑₐ₂ᶿ?

The energy required for 1 mol of gaseous 1- ions to acquire an electron to form 1 mol of gaseous 2- ions. \[***Endothermic***\]

[5.2.3] What is meant by the term ‘oxidising agent?

A species that oxidises another species, and is itself reduced.

[5.2.3] What is meant by the term ‘reducing agent?

A species that reduces another species, and is itself oxidised.

[5.2.3] What are the half-equation rules?

• Balance all atoms that are not O or H.

• Balance O by adding H₂O.

• Balance H by adding H⁺.

• Balance Charge by adding e⁻.

• Add State Symbols.

[5.2.3] What is meant by the standard electrode potential of a cell?

The E° of a half-cell is the electromotive force (e.m.f) of a half cell compared with a standard hydrogen half cell measured at standard conditions.

The more negative the E°, the larger the tendency for the half cell to be oxidised (to release electrons).

[5.2.3] What value are standard hydrogen electrodes assigned?

0.00v

[5.2.3] What does a standard hydrogen electrode look like?

[5.2.3] Why are platinum electrodes typically used?

They are inert and conductive.

[5.2.3] What are the standard conditions for electrode potentials?

298K, 100kPa, 1.00moldm⁻³.

[5.2.3] What does an electrochemical cell look like?

[5.2.3] What are the three types of half cell?

• Metals in contact with solutions of their ions.

• Gases in contact with solutions of their ions.

• Solutions of Ions in two different oxidation states.

[5.2.3] How can the cell potential be calculated?

Positive Value - Negative Value

[5.2.3] What are the three types of cell?

• Non-Rechargeable Storage Cell

• Rechargeable Storage Cell

• Fuel Cell

[5.2.3] What is a non-rechargeable storage cell?

A cell that is discarded once the voltage falls.

[5.2.3] What is a rechargeable storage cell?

A cell whereby the reaction is reversed, and its chemicals are regenerated.

[5.2.3] What is a fuel cell?

A cell that generates electrical energy using external supplies of a fuel and oxidising agent.

[5.2.3] What are fuel cells made up of?

Anode, Cathode, Electrolytes.

[5.2.3] What is the net result of the occurring reactions in a fuel cell?

• Fuel reacts with oxygen.

• Water is produced.

  • Voltage is produced.

[5.2.3] What does a typical hydrogen-oxygen fuel cell look like?

This specific example uses an acidic electrolyte.

[5.2.3] What is an alkaline hydrogen-oxygen fuel cell?

Alkaline fuel cells use alkaline electrolytes and the membrane only allows for the movement of hydroxide ions.

[5.3.1] What are transition metals?

D-block elements that form at least one ion with an incomplete d-subshell.

[5.3.1]

[5.3.1] What shape is formed by [Fe(NH₃)₄(H₂O)₂]²⁺?

Octahedral (bond angle of 90 degrees).

[5.3.1] What shape is formed by [CuCl₄]²⁻?

Tetrahedral (bond angle of 109.5 degrees).

[5.3.1] Why are tetrahedral shapes formed by transition metals?

When the ligands are larger.

[5.3.1] Which metals form square planar molecules?

Group 10 Molecules, e.g. Ni/Pt

[5.3.1] What shape is formed by AgCl₂?

Linear (bond angle of 180 degrees).

[5.3.1] What are the cis/trans isomers of [Pd(NH₃)₂Cl₂]?

• In the cis-isomer, the bond angle is 90°.

• In the trans-isomer, the bond angle is 180°.

[5.3.1] What is the structure of cis-platin?

[5.3.1] How does cis-platin work as an anti-cancer drug?

• Loses chloride ligands.

• Binds to DNA, causing it to bend.

• Inhibits DNA synthesis.

[5.3.1] How does [Cu(NH₂CH₂CH₂NH₂)₂(H₂O)₂]²⁺ present cis/trans isomerism?

[5.3.1] How does cis-[Cu(NH₂CH₂CH₂NH₂)₂(H₂O)₂]²⁺ present optical isomerism?

[5.3.1] Why are carbons present in bidentate ligands?

In order to allow them to ‘stretch’ to bond to two parts of the octahedral shape.

[5.3.1] What are ligand substitution reactions?

A reaction in which one ligand is replaced by another.

[5.3.1] How is [Cu(NH₃)₄(H₂O)₂]²⁺ formed from [Cu(H₂O)₆]²⁺?

[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O

[5.3.1] What is the colour change associated with the formation of [Cu(NH₃)₄(H₂O)₂]²⁺ from [Cu(H₂O)₆]²⁺?

Blue —> Dark Blue

[5.3.1] How is [CuCl₄]²⁻ formed from [Cu(H₂O)₆]²⁺?

[Cu(H₂O)₆]²⁺ + 4HCl ⇌ [CuCl₄]²⁻ + 6H₂O + 4H⁺

[5.3.1] What is the colour change associated with the formation of [CuCl₄]²⁻ from [Cu(H₂O)₆]²⁺?

[CuCl₄]²⁻ is a yellow product.

• The resulting solution is green as there is an equilibrium between [Cu(H₂O)₆]²⁺ (blue) and [CuCl₄]²⁻ (yellow).

[5.3.1] How is [Cr(NH₃)₆]³⁺ formed from [Cr(H₂O)₆]³⁺?

[Cr(H₂O)₆]³⁺ + 6NH₃ ⇌ [Cr(NH₃)₆]³⁺ + 6H₂O

[5.3.1] What is the colour change associated with the formation of [Cr(NH₃)₆]³⁺ from [Cr(H₂O)₆]³⁺?

Violet —> Purple

[5.3.1] Why is Fe²⁺ important in haemoglobin?

Binds to oxygen to form oxyhaemoglobin, in order to transport it around the body.

[5.3.1] Why is the ligand substitution in Fe²⁺ of oxygen to carbon monoxide bad?

CO binds to Hb more strongly than O₂, which prevents O₂ being transported around the body.

• The equilibrium makes it nearly impossible for CO to be substituted for another ligand, meaning that if carboxyhaemoglobin concentration becomes too high, death can occur.

[5.3.1] What is a precipitation reaction?

When two aqueous solutions containing ions react together to produce a precipitate.

[5.3.1] How is Cu(OH₂) formed from [Cu(H₂O)₆]²⁺?

Cu²⁺₍ₐᵩ₎ + 2NaOH₍ₐᵩ₎ —> Cu(OH)₂₍ₛ₎

• Ionic Equation = [Cu(H₂O)₆]²⁺₍ₐᵩ₎ + 2OH⁻₍ₐᵩ₎ —> [Cu(H₂O)₄(OH)₂]₍ₛ₎ + 2H₂O₍ₗ₎

[5.3.1] What is the colour change in the formation of Cu(OH)₂ from Cu²⁺?

Blue —> Light Blue

[5.3.1] How is Fe(OH)₂ formed from Fe²⁺?

Fe²⁺₍ₐᵩ₎ + 2OH⁻₍ₐᵩ₎ —> Fe(OH)₂₍ₛ₎

[5.3.1] What is the colour change in the formation of Fe(OH)₂ from Fe²⁺?

Pale Green —> Dark Green

[5.3.1] How is Fe(OH)₃ formed from Fe³⁺?

Fe³⁺₍ₐᵩ₎ + 3OH⁻₍ₐᵩ₎ —> Fe(OH)₃₍ₛ₎

[5.3.1] What is the colour change in the formation of Fe(OH)₃ from Fe³⁺?

Yellow —> Red/Brown

[5.3.1] How is Cr(OH)₃ formed from Cr³⁺?

Cr³⁺₍ₐᵩ₎ + 3OH⁻₍ₐᵩ₎ —> Cr(OH)₃₍ₛ₎

[5.3.1] What is the colour change in the formation of Cr(OH)₃ from Cr³⁺?

Purple —> Green/Grey

[5.3.1] How is [Cr(OH)₆]³⁻ formed from [Cr(H₂O)₆]³⁺?

[Cr(H₂O)₆]³₍ₐᵩ₎+ 6OH⁻₍ₐᵩ₎ —> [Cr(H₂O)₆]³⁺₍ₛ₎ + 6H₂O

[5.3.1] What is the colour change in the formation of [Cr(OH)₆]³⁻ from [Cr(H₂O)₆]³⁺?

Purple —> Dark Green

[5.3.1] How is Mn(OH)₂ formed from Mn²⁺?

Mn²⁺₍ₐᵩ₎ + 2OH⁻₍ₐᵩ₎ —> Mn(OH)₂₍ₛ₎

[5.3.1] What is the colour change in the formation of Mn(OH)₂ from Mn²⁺?

Pale Pink —> Buff

[5.3.1] What happens when you add excess NH₃ to [Cu(H₂O)₆]²⁺?

• It will initially form a Cu(OH)₂ precipitate.

  • Eventually it becomes a ligand substitution reaction, and becomes [Cu(NH₃)₄(H₂O)₂]²⁺.