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A2 Chemistry > Unit 5 > Flashcards

Unit 5 Flashcards

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1
Q

Enthalpy of sublimation

A

Enthalpy change for a solid metal turning into gaseous atoms. Same as atomisation values.

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2
Q

Bond dissociation Enthalpy

A

Standard molar Enthalpy change when 1 mole of a covalent bond is broken into 2 gaseous atoms (free radicals).
For diatomic, ΔH diss is same as 2x ΔH at.

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3
Q

First electron affinity

A

Enthalpy change when 1 mole of gaseous atoms gains 1 mole of electrons to form 1 mole of gaseous -1 ions.
Exo for atoms than normally form -ve ions (more stable) and attraction between nucleus and electron.

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4
Q

Second electron affinity

A

Enthalpy change when 1 mole of gaseous -1 ions gains 1 electron per ion to form gaseous -2 ions.
Endo as energy is required overcome repulsion of -ve ion and electron.

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5
Q

Enthalpy of lattice formation

A

Standard Enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form.

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6
Q

Enthalpy of lattice dissociation

A

Standard Enthalpy change when 1 mole of an ionic crystal lattice is separated into its constituent ions in gaseous form.

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7
Q

Enthalpy of hydration

A

Enthalpy change when 1 mole of gaseous ions become aqueous ions (always exo).

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8
Q

Enthalpy of solution

A

Standard Enthalpy changed when 1 mole of an ionic solid dissolves in a large amount of water, so that the dissolved ions are separated and don’t interact with each other.

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9
Q

Enthalpy of atomisation

A

Enthalpy change of when 1 mole of gaseous atoms is formed from the element in its standard state.

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10
Q

Born haber cycles

A

Indirectly calculating lattice Enthalpy using available data and link together in an Enthalpy cycle.

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11
Q

Strength of Enthalpy of lattice formation depends on:

A

Size of ions

Charge on ion.

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12
Q

Perfect ionic model

A

Ions of lattice Enthalpies are:
100% ionic
Spherical
Purely electrostatic attractions

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13
Q

There is a tendency to covalent character in ionic substances when:

A

+ve ion is small
+ve ion has multiple charges
-ve ion is large
-ve ion has multiple charges

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14
Q

Spontaneous process

A

Proceeds on its own without external factors.
Mainly exo reactions as they have products more thermodynamically stable.
Some are endo.

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15
Q

Entropy

A

Description of no. of ways atoms can share quanta of energy (no. of ways of arranging atoms and energy)
The more ways, the more disordered, the higher the entropy.

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16
Q

Increase in entropy when:

A

Change in state from solid/liquid to gas.

Significant increase in no. of molecules between products and reactants.

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17
Q

Equation for entropy change, ΔS

A

ΔS= ΣS products - ΣS reactants

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18
Q

Gibbs free energy change, ΔG

A

Balance between entropy and Enthalpy. Determines feasibility of a reaction.

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19
Q

Equation for free energy change, ΔG

A

ΔG= ΔH - TΔS

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20
Q

Units for calculating free energy change

A

T: K
S: J k^-1 mol^-1
ΔG: KJ mol^-1
ΔH: KJ mol^-1

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21
Q

Equation for Enthalpy of solution, ΔH sol

A

ΔH sol = ΔH Ldiss + ΣΔH hyd
OR
ΔH sol = ΣΔH hyd - ΔH Lform

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22
Q

Mean bond energy

A

Enthalpy required to break a covalent bond into gaseous atoms, averaged over different molecules.

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23
Q

Equation for mean bond energy

A

ΔH = Σbond energies broken - Σbond energies made

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24
Q

Reactions of Na and Mg with H2O

A

Na reacts with cold water, fizzes on surface.

Mg reacts slowly with cold water, but reacts with steam to form an oxide.

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25
Q

Trends of oxides of elements

Na –> S

A

Ionic metal oxides–> basic behaviour
Non metal covalent oxides—> acidic behaviour

Al2SO3 (high strength of ionic lattice) and SiO2 (macromolecular structure) don’t dissolve in water. pH 7

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26
Q

Amphoteric

A

Can react and dissolve as a base and an acid.

E.g. Al2O3

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27
Q

Electrochemical cells

A

Has 2 half cells.
Connected by a salt bridge.
Simple half cells have a metal (electrode) and a sol. of compound metal.
Half cells will produce small voltage if connected to a circuit.

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28
Q

Potential difference is measured with…

A

a high resistance voltmeter.

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29
Q

Why use a high resistance voltmeter

A

high resistance to stop current from flowing ∴ reactions will not occur.
can measure a possible max. E.

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30
Q

Salt bridge

A

used to connect up circuit, free moving ions conduct charge.
made from filter paper soaked in salt sol. (potassium nitrate).
salt should be unreactive with electrodes and electrode sol.
wire not used as will form its own electrode system.

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31
Q

When current flows…

A

+ve electrode will undergo reduction.
-ve electrode will undergo oxidation.
voltage falls to 0 as reactants are used up.

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32
Q

Cell diagrams

A 
solid line= boundary between phases
double line= salt bridge
voltage indicated
the more +ve half cell is on the right. 
most oxidised form is next to double line.
commas are used to separate oxidised from reduced species.
solid line separates physical states
H+ and H2O can be left out.
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33
Q

Systems that don’t include metals

A

use a platinum electrode (included in cell diagram):
provides conducting surface for e- transfer
unreactive
can conduct electricity

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34
Q

Standard Hydrogen Electrode (SHE)

A

potential of all electrodes are measured by comparing their potential to SHE.
SHE potential is 0.
represented by: Pt| H2| H+

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35
Q

Components of SHE

A

H2 gas at 100KPa
Sol. containing H ions at 1M conc.
Temp. 298K

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36
Q

Secondary standards (for SHE)

A

easier standards are used for SHE, they are calibrated against SHE.
e.g. silver/ silver chloride
calomel electrode

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37
Q

Standard electrode potentials (SEP)

A

potential difference measured when an electrode system is connected to SHE, standard cond.s apply.

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38
Q

Standard cond.s for SEP

A

all ion sol.s at 1M conc.
temp 298K
100KPa pressure
No current flowing.

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39
Q

Equation for EMF

A

EMF = Erhs - Elhs (right - left)

40
Q

Using electrode potentials

A

More -ve half cell will always oxidise and go backwards, e- goes to +ve electrode.

More +ve half cell will always reduce and go forward.

41
Q

Using series of SEP

A

More +ve has increasing tendency for species on left to reduce, and act as oxidising agent (higher ox. no.)

More -ve has increasing tendency for species on right to oxidise, and act as reducing agent (lower ox. no.)

42
Q

Equation for Ecell from 2 SEPs

A

Ecell = Ered - Eox ( reduction - oxidation)

43
Q

Ecell

A

A measure of how far from equilibrium cell reaction is.

The more +ve the Ecell, the more likely the reaction is going to occur.

44
Q

Effect of conc. on Ecell

A

Increasing conc. of reactants –> increase Ecell

Decrease conc. of reactants –> decrease Ecell

45
Q

Effect of temp on Ecell

A

As most cells are exo in spontaneous direction, increasing temp decreases Ecell, as equilibrium shifts backwards.
If Ecell is +ve, reaction might occur.
Reaction will not occur with high Ea.

46
Q

Cells

A

can be rechargeable, non-rechargeable (irreversible) and fuel cells.
Reversible cells only work if products stay attached to electrode and doesn’t disperse.

47
Q

Fuel cells

A

Uses energy from reaction of a fuel with O2 to create voltage.

48
Q

Standard cond.s for fuel cells

A

Rate is too slow, ∴higher temp is used, as reaction is exo, Ecell falls.
∴Higher pressure
continuously fed O2 and H2 –> maintains constant voltage.

49
Q

Advs of Fuel cells

A

less pollution + CO2

greater efficiency.

50
Q

Disadvs of Fuel cells

A 
Expensive
Storing and transporting
Feasibility of pressured liquid
Limited life cycle of a solid adsorber or absorber.
Limited lifetime
Use of toxic chemicals in production.
51
Q

Ethanol fuel cells

A

Can be made from renewable sources, in carbon neutral way.
Abundant raw materials to produce ethanol by fermentation.
Less explosive than H2.

52
Q

Transition metals properties

A 
from elements Sc-->Cu
have incomplete d sub shells in either atoms or ions
complex formation
form coloured compounds
variable oxidation states
used as catalysts
53
Q

Complex

A

central metal ion surrounded by ligands.

54
Q

Ligands

A

An atom, ion or molecule which donates a lone pair of electrons.
forming dative bonds to form a complex.

55
Q

Coordination no.

A

No. of co ordinate bonds formed to a central metal ion.

56
Q

Unidentate

A

forms 1 dative bond per ligand. e.g. H2O, NH3, Cl-

57
Q

Bidentate

A

has 2 atoms with lone pairs, can form 2 dative bonds per ligand. e.g. NH2CH2CH2NH2, C2O4^2-

58
Q

Multidentate

A

Forms 3 or more dative bonds per ligand. e.g. EDTA^4- (forms 6 dative bonds per ligand).

59
Q

Things that cause colour change

A

Oxidation State
Coordination no.
Ligand

60
Q

How colour arises

A

Electronic transitions from ground state to excited state: difference between d orbitals.
Part of visible light is absorbed so d electrons go to higher energy levels.
The light that is not absorbed is transmitted to give the substance colour.

61
Q

Equation for energy difference

A

ΔE = hv
ΔE - Energy difference between split orbitals.
h - Planck’s constant, 6.63x10^-34Js
v - freq. of light absorbed
ligands cause 5d orbitals to split into 2 energy levels.

62
Q

Compounds without colour

A

when haven’t got any d electrons to move (Sc^3+)
or when have full d subshell and no space for electron transfer (Zn^2+ and Cu^+).
Both mean no energy transfer equal to visible light.

63
Q

Spectrophotometry

A

Some light is absorbed when visible light of increasing freq. is passed through a sample of coloured complex ion.
Amount of light absorbed ∝ conc. of absorbing species.
∴ Absorption o visible light is used to determine conc. of coloured ions.

64
Q

Process for Spectrophotometry

A
  1. Add ligand
  2. Make up sol. of known conc.
  3. Measure absorption or transmission.
  4. Plot graph of results or calibration curve.
  5. Measure absorption of unknown and compare.
65
Q

Variable oxidation states

A

Stability of 2+ with respect to 3+ increases across period.
Compounds with high Ox. states –> oxidising agents. e.g. MnO4^-
Compounds with low Ox. states –> reducing agents. e.g. V^2+, Fe^2+

66
Q

Reducing Chromium

A

Cr^3+ (green) and Cr^2+ (blue) are formed by reduction of Dichromate (orange) by strong reducing agent: Zn in HCL sol.
Fe^2+ will only reduce to Cr^3+ (not strong enough).
Reduction with Fe^2+ can be used as a quantitative redox titration (doesn’t need an indicator).
Zn and dichromate needs to be under a hydrogen atmosphere.

67
Q

Manganate redox titration

A

Between Fe^2+ and MnO4^- (purple)
Self indicating because of significant colour change from reactant to product.
Colourless –> purple
If Mn in burette then end point is at first permanent pink colour.

68
Q

Choosing correct acid for Manganate titration

A

Acid is needed to supply 8H+ ions
Some acids set up alternative redox reactions ∴ make it inaccurate.
Only use dilute H2SO4
Too little vol. of H2SO4 and using a weak acid (like ethanoic acid) means it can’t supply enough H+ ions, ∴ MnO2 is produced instead.
Its brown colour will mask colour change ∴ more vol. of Mn used ∴ inaccurate.
Cant use conc. HCL (Cl- ions oxidise) or Nitric acid (oxidising agent) either.

69
Q

Oxidation in alkaline solutions

A

When TMs in low Ox. states are in in alkaline sol. –> more easily oxidised than when in acidic sol.
easier to remove e- from -ve ion.
Oxidising agents such as H2O2 used.

70
Q

Oxidation of chromium in alkaline sol.

A

Oxidising agent = H2O2

Green sol. –> yellow sol.

71
Q

Oxidation of cobalt in alkaline sol.

A

Oxidising agent = H2O2

blue ppt –> brown ppt

72
Q

Half equations in alkaline cond.s

A

Balance as if in acidic cond.s, then add OH- to convert to alkaline.

  1. Add H2O to balance O
  2. Add H+ to balance H
  3. Add OH- to both sides to cancel H+
73
Q

Ammoniacal Oxidation of cobalt

A

Ammonia ligands make Co(II) state unstable + easier to oxidise.
Air oxidises Co(II) to Co(III).

74
Q

Chromate/ dichromate equilibrium

A

they can be converted from 1 to the other.
yellow sol. (Chromate) –> orange sol. (Dichromate)
Acid base reaction
adding acid –> equilibrium shifts to dichromate
adding alkali –> equilibrium shifts to chromate as removes H+ ions.

75
Q

Heterogeneous catalyst

A

In different phase from reactants.
usually solid (reactants are gas or in sol.)
higher conc. of reactants at solid surface, ∴ higher collision freq.
TMs can use 3d and 4s e- of atoms on metal surface to form weak bonds to reactants.

76
Q

Homogeneous catalyst

A

In same phase as reactants.
Reaction proceeds through intermediate species.
which has different Ox. state to TM.
At end of reaction original Ox. state reoccurs.

77
Q

Strength of adsorption of heterogeneous catalysts

A

some metals have too strong adsorption –> products cant be released.
some metals have too weak adsorption –> reactants don’t adsorb in high enough conc.
Ni and Pt have correct strength

78
Q

Steps in heterogeneous catalysis

A
  1. Reactants form bonds with atoms on active sites of surface of catalyst (adsorption of reactants).
  2. ∴ bonds in reactants are weakened and break.
  3. New bonds form, held close together on catalyst surface.
  4. Weakens bonds between product and catalyst and product leaves (desorbs).
79
Q

Examples of heterogeneous catalysts

A
  1. V2O5 used as a catalyst in Contact Process.
  2. Cr2O3 catalyst used in manufacture of methanol from CO and H2.
  3. Fe catalyst in Haber Process.
80
Q

Examples of homogeneous catalysts

A
  1. Reaction between iodide and persulphate ions, catalysed by Fe^2+, makes collision between oppositely charged ions, ∴ lower Ea.
  2. Autocatalytic reaction between ethanedioate and manganate ions, Mn^2+ product is the autocatalyst. initial reaction is slow.
    Follow reaction rate using a spectrometer, or titrating samples.
81
Q

Constructing catalysed mechanism for a reaction

A
  1. Split full equation into its 2 half equations
  2. Add catalyst to make 2 new redox equations.
  3. Make sure oxidised equation is combined with original reduced half equation and vice versa.
  4. Check 2 mechanism equations add up to original full non-catalysed equation.
82
Q

Other applications of TM complexes

A
  1. Fe(II) in haemoglobin enables O2 to be transported in the blood.
  2. Pt(II) complex cisplatin is used as anticancer drug.
83
Q

Silver chemistry

A

Ag+ commonly forms linear complexes
Like TM in the way it can form complexes and act as catalyst.
But, all complexes have full 4d sub shell and 1+ oxidation state, ∴ not TM.
Doesn’t form coloured compounds or have a variable oxidation state, as not able to do electron transitions between d orbitals –> no coloured compounds.

84
Q

Reactions of halides with silver nitrate

A

Fluoride –> No ppt
Chloride –> White ppt
Bromide –> cream ppt
Iodide –> pale yellow ppt

85
Q

To help distinguish between silver halides, ammonia sol. is added…

A

AgCl –> dissolves in dilute ammonia to give complex ion (colourless)
AgBr –> dissolves in conc. ammonia to give complex ion (colourless)
AgI –> doesn’t react with ammonia - too insoluble.
[Ag(NH3)2]^+ is used in Tollen’s reagent.

86
Q

Autocatalyst

A

Where one product can catalyse the reaction.

87
Q

Metal aqua ions

A

Metal ions formed in Aq sol.
2+ ions (Fe-green, Co-pink, Cu-blue)
3+ ions (Al-colourless, Cr- ruby/ green, Fe-violet)

88
Q

Acidity or hydrolysis reactions

A

Equilibria happens in aq sol. of metal ions.
Leads to generation of very acidic solutions of M^3+ ions and very weakly acidic solutions with M^2+ ions.
greater acidity in terms of polarising power of 3+ ions.
the greater the polarising power, the more strongly it attracts H2O molecule, weakens O-H bond, so it breaks more easily.

89
Q

Reactions with limited OH- and limited NH3

A

Bases OH- and ammonia, when limited, form the same hydroxide ppts.
Form in Deprotonation acid base reactions.
2+ ions (Cu-blue ppt, Co-blue ppt, Fe(II)-green ppt)
3+ ions (Cr(III)-green ppt, Fe(III)-brown ppt, Al-white ppt)

90
Q

Reactions with excess OH-

A

With excess NaOH, Cr and Al hydroxides dissolve, they are Amphoteric.
Cr –> green sol.
Al –> colourless sol.

91
Q

Reactions with excess NH3

A

ligand substitution occurs with Cu, Co and Cr, their ppts dissolve, without a change of coordination no. for Co and Cr. (sub with this may be incomplete, as with Cu)
Ligands- NH3 and H2O are similar in size and are uncharged.
Cr –> purple sol.
Co –> pale yellow sol.
Cu –> deep blue sol.
NH3 is acting as lewis base.

92
Q

Reactions with chloride ions

A

high conc. Cl- ions with Aq ions leads to ligand sub reaction.
Cl- ligand is larger than uncharged H2O and NH3 ligands ∴ ligand exchange can involve change in coordination no.
Cu –> yellow/ green sol.
Co –> blue sol.

93
Q

Reactions with carbonate solutions

A

2+ ions –> MCO3 ppt (Cu-blue/green, Co-pink, Fe(II)-green)
3+ ions –> M(OH)3 ppt and CO2 (Al-white, Cr(III)-green, Fe(III)-brown)
M(OH)3 is formed due to the higher polarising power of 3+ ion.

94
Q

Stability of complexes

A

sub of unidentate with bidentate or multidentate leads to more stable complex.
As +ve entropy change in these reactions, due to more molecules of products than reactants.

95
Q

Common ligands

A
  1. Ethanedioate
  2. Ethane-1,2-diamine: bidentate. ΔG will be -ve, as ΔS is +ve and ΔH is close to 0, due to no. of dative bonds and type are same ∴ energy needed to make/ break bonds is same. This as a base, can carry out deprotonation reactions forming hydroxide ppts.
96
Q

Quantitative calculations with complex ions

A

E.g. EDTA titrations

1: 1 ratio with any metal ion
1. Find moles
2. Use balanced equation to find moles of ion.
3. Find conc. of ion in specific vol.

97
Q

using silver nitrate to work out formula of chloride containing complexes

A

Sometimes compounds have complex with Cl- ions inside (ligands) and outside complex (attracted ionically)
If silver nitrate is added, only form silver chloride ppt with free Cl- ions outside complex.

e. g.1. Co(NH3)6Cl3, 1:3 mole ratio with silver nitrate, as 3 free Cl- ions outside complex.
e. g.2. Cr(NH3)5Cl3, 1:2 ratio with silver nitrate, as 2 free Cl- ions outside, 1 ligand (inside).
e. g.3. Cr(NH3)4Cl3, 1:1 ratio with silver nitrate, as 1 free Cl- ion outside, 2 ligands.

A2 Chemistry (2 decks)

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