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Revise Deck Flashcards

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1
Q

Suggest how the ligand of EDTA⁴⁻ can be formed from EDTA.

A

Placing EDTA in an alkaline buffer solution, causing the acid COOH groups to lose their H.

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2
Q

What is required for optical isomerism in octahedral complexes?

A

Bidentate ligands.

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3
Q

Why may spectroscopy be used instead of colorimetry?

A

It uses wavelengths in the UV region as well as visible light.

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3
Q

What is the general colour of:
copper(II) compounds
iron(II) compounds
iron(III) compounds
cobalt(II) compounds

A

blue
green
brown
pink

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4
Q

What is the equation that links the energy difference between the ground and excited state and frequency of the wavelength?

A

∆E = hv
∆E is difference in energy (J), h is Planck constant (given, Js), v is frequency (Hz or s⁻¹)

c = vλ

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5
Q

Explain what happens when hexaaqua copper(II) reacts with:
sodium hydroxide
ammonia solution
sodium carbonate solution

A

NH₃: blue ppt and acid-base. When in excess NH₃, deep blue solution forms:
Cu(H₂O)₄(OH)₂ + 4NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ + 2H₂O + 2OH⁻
This is ligand substitution.

Na₂CO₃: CuCO₃ green ppt forms.

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6
Q

Explain what happens when hexaaqua iron(II) reacts with:
sodium hydroxide
ammonia solution
sodium carbonate solution

A

neutral complex. No change when excess NaOH.

NH₃: green ppt and acid-base. No change when excess NH3.

Na₂CO₃: FeCO₃ green ppt forms.

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6
Q

Explain what happens when hexaaqua iron(III) reacts with:
sodium hydroxide
ammonia solution
sodium carbonate solution

A

Originally yellow/brown solution.

NaOH: brown ppt forms; three hydroxide ions react, causing an acid-base reaction to form a neutral complex. No change when excess NaOH.

NH₃: brown ppt and acid-base. No change when excess NH3.

Na₂CO₃: Fe(OH)₃(H₂O)₃ brown ppt produced (acid-base) and bubbles of gas are produced from CO₂ released.

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7
Q

aluminium(III) reacts with:
sodium hydroxide
ammonia solution
sodium carbonate solution

State any observations.

A

colourless solution

NaOH: white ppt forms; three hydroxide ions react, causing an acid-base reaction to form a neutral complex. Redissolves to form colourless solution in excess NaOH: [Al(OH)₆]³⁻ /[Al(OH)4]-

NH₃: white ppt and acid-base. No change when excess NH3.

Na₂CO₃: Al(OH)₃(H₂O)₃ white ppt produced (acid-base) and bubbles of gas are produced from CO₂ released.

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8
Q

State the colours and oxidation states of vanadium of the following ions in solution:
VO₂⁺
VO²⁺
V³⁺
V2⁺

A

You better Get Vanadium
+5, yellow
+4, blue
+3, green
+2, violet

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9
Q

How can a vanadium species be reduced?

A

Using a reducing agent (zinc) in the presence of acid (HCl).

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10
Q

Write the formula of the manganate ion, stating its colour in solution and oxidation state of manganese.

A

MnO₄⁻, purple solution, +7

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11
Q

Explain why a support medium may be used for a heterogenous catalyst.

A

The use of a support medium maximises the surface area of a heterogeneous catalyst and so minimises the cost.

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12
Q

Explain, with the aid of equations, how V2O5 acts as a catalyst in the Contact process.

A

SO₂ + V₂O₅ → SO₃ + V₂O₄
V₂O₄ + 0.5O₂ → V₂O₅

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13
Q

Explain, with the aid of equations, how Fe2+ ions catalyse the reaction between I− and S2O82-.

A

2Fe²⁺ + S₂O₈²⁻ → 2Fe³⁺ + 2SO₄²⁻
2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂

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14
Q

Explain, with the aid of equations, how Mn2+ ions autocatalyse the reaction between C2O42- and MnO4-.

A

5C₂O₄²⁻ + 2MnO₄⁻ + 16H⁺ → 10CO₂ + 2Mn²⁺ + 8H₂O

4Mn²⁺ + MnO₄⁻ + 8H⁺ → 5Mn³⁺ + 4H₂O
2Mn³⁺ + C₂O₄²⁻ → 2Mn²⁺ + 2CO₂

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15
Q

Why does Fe III not produce a carbonate when reacting with a metal carbonate?

A

Acidic (high enough charge density), polarises the H20 molecule, 3 H20 deprotonate

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16
Q

Define standard electrode potential (E°)

A

electrode potential of a standard electrode with ion concentration 1.00moldm⁻³ at 298K connected to a standard hydrogen electrode with H⁺ ion concentration of 1.00moldm⁻³, with 100kPa H₂ gas at 298K, using a high-resistance voltmeter and a salt bridge.

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17
Q

conditions for SHE

A

1.00moldm⁻³, with 100kPa H₂ gas at 298K, Pt (black)

18
Q

What is the difference between a primary cell and a secondary cell?

A

A primary cell is non-rechargeable (irreversible), whereas a secondary cell is rechargeable.

19
Q

Define pH.

A

pH = -log₁₀[H⁺]

20
Q

Define Kw, stating its units.

A

Kw is the ionic product of water.
Kw = [H⁺][OH⁻]
mol²dm⁻⁶

21
Q

What does a higher Ka indicate?

A

A ‘stronger’ weak acid.

22
Q

What does a higher pKa value indicate?

A

A ‘weaker’ weak acid.

23
Q

Write the Henderson-Hasselbalch equation.

A

pH = pKa + log([salt]/[acid])

24
Q

Define enthalpy change.

A

The heat energy change of a reaction at constant pressure.

25
Q

Define enthalpy of lattice formation.

A

The enthalpy change when one mole of an ionic compound is formed from its constituent gaseous ions.
Exothermic

26
Q

Define standard enthalpy of formation. (∆fH⦵)

A

The enthalpy change when one mole of a compound is formed from its constituent elements, when all products and reactants are in their standard states under standard conditions.

27
Q

Define enthalpy of atomisation. (∆atH⦵)

A

enthalpy when 1 mole of gaseous atoms is formed from the element in standard cond.

28
Q

Write an equation, including state symbols, for the process that has an enthalpy change equal to the enthalpy of atomisation of:

sodium

A

Na(s) → Na(g)

29
Q

Define first electron affinity. (∆EAH⦵)

A

The enthalpy change when one mole of gaseous atoms forms one mole of gaseous 1- ions.

30
Q

Explain what is shown if the experimental value for lattice enthalpy of a compound is:
∙ higher than the theoretical value.
∙ close to the theoretical value.

A

Explain what is shown if the experimental value for lattice enthalpy of a compound is:
∙ higher than the theoretical value.
∙ close to the theoretical value.

The compound is ionic with some covalent character, which adds additional strength to the bond (requiring more energy to break).

The compound is ionic.

31
Q

Define enthalpy of hydration.

A

The enthalpy change when one mole of gaseous ions is converted to one mole of aqueous ions.

32
Q

Write the equation for Gibbs free energy.
State the units for each term.

A

∆G⦵ = ∆H⦵ - T∆S⦵

G and H are kJ mol⁻¹
T is K
S is J K⁻¹mol⁻¹

33
Q

What is a Maxwell-Boltzmann distribution?

A

A plot of the number of gaseous molecules against the energy they have at a fixed temperature.

34
Q

Explain what will happen to the Maxwell-Boltzmann distribution if a catalyst is added to the reaction mixture. How does this affect rate?

A

Explain what will happen to the Maxwell-Boltzmann distribution if a catalyst is added to the reaction mixture. How does this affect rate?

Ea will shift to the left relative to the x axis, since catalysts provide an alternative reaction pathway with a lower activation energy.
The shaded area under the curve will now be larger,

35
Q

Describe how temperature affects the sketch of a MBD

A

Sketch a higher temperature and a lower temperature on the same molecular energies distribution. Describe how temperature affects the sketch.

Lower temperature distributions are moved to the left and the peak is higher.
Higher temperature distributions are moved to the right and the peak is lower.
All curves need to start at the origin.

36
Q

Describe how concentration affects the sketch of a MBD.

A

The curve retains the same basic shape. The most probable energy (and average energy) remains the same, so the peak will be higher at the same energy.

37
Q

State the conditions for the Haber process.
State any compromises.

A

State the conditions for the Haber process.
State any compromises.

450⁰C - between rate and yield
200 atm - between yield and cost
iron catalyst

38
Q

Explain the trend in melting point down group 2.

A

Explain the trend in melting point down group 2.

Generally decreases, since atomic radius increases therefore the electrostatic attraction between the positive metal ions and the delocalised electrons decrease. Therefore, the strength of the metallic bonds decrease which requires less energy to break in order to melt the metal.

39
Q

Explain a use of magnesium.

A

Explain a use of magnesium.

Extract titanium from titanium chloride:
TiCl4 + 2Mg –> 2MgCl2 + Ti

40
Q

Describe a test to identify the presence of sulphate ions.

A

Add HCl
Add BaCl2
A thick white precipitate forms (BaSO4).

41
Q

State the trends of the halogens down the group. (atomic r, reactivity, bp, electroneg)

A

The colour of the elements becomes darker.
Atomic radius increases.
Reactivity decreases
Boiling points increase.
Electronegativity decreases.

42
Q

Give the ionic equation and symbol equation for the reaction of cold, dilute, aqueous sodium hydroxide with chlorine.
State an observation.
What type of reaction is this?

A

Give the ionic equation and symbol equation for the reaction of cold, dilute, aqueous sodium hydroxide with chlorine.
State an observation.
What type of reaction is this?

Cl2 + 2OH- –> Cl- + ClO- + H2O
Cl2 + 2NaOH –> NaCl + NaClO + H2O

Yellow-green gas forms a colourless solution.
Redox reaction - disproportionation

43
Q

Give the reversible reaction for chlorine and water. Name the products formed and the oxidation state of chlorine in each.

and state the reaction occuring in bright light

A

Cl2 + H2O ⇌ HCl + HClO
Hydrochloric acid - -1
Hypochlorous acid / chloric (I) acid - +1

2Cl2 + 2H2O –> 4HCl + O2
UV light breaks down the HClO into HCl and O2.

44
Q

Give the equation for the reactions of:
- sodium fluoride and concentrated sulphuric acid
- sodium chloride and concentrated sulphuric acid.

A

NaF + H2SO4 –> NaHSO4 + HF
NaCl + H2SO4 –> NaHSO4 + HCl

45
Q
A

Chemistry (33 decks)

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