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developing metals Flashcards

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1
Q

use of redox titrations

A

used to find the amount of oxidising agent needed to react with a known concentration of reducing agent or vice versa

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2
Q

manganate(VII) titration method

A

1) Add a known mass of impure Fe2+(aq) to a conical flask
2) Add dilute sulfuric acid in excess (because the equation of the reaction contains H+ ions)
3) Add MnO4-(aq) to the flask with the burette until the solution in the conical flask turns slightly pink and remains this colour
4) Repeat until you have two concordant results

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3
Q

coordination number meaning

A

the number of coordinate bonds from ligands to the central atom or ion

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4
Q

shape and angle of molecule with coordination number of 6

A

octahedral shape, 90 degrees
because
- 6 lone pairs donated to the central metal
- move as far apart as possible to minimise repulsion, hence 90 degrees

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5
Q

shape and angle of molecule with coordination number of 4

A

tetrahedral shape (109.5 degrees) or square planar (90 degrees)
because
- 4 lone pairs donated to the central metal
- move as far apart as possible to minimise repulsion, hence 109.5 / 90

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6
Q

finding coordination number

A

(if monodentate) coordination number can be found by looking at the small number of the ligand formula

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7
Q

oxidation state rules

A

elements = 0
simple ions = the charge on the ion
group 1 = +1
group 2 = +2
Al = +3
F = -1
H = +1, unless metal hydrides
O = -2, unless OF2, peroxides, superoxides
Cl = -1, unless oxygen compounds
sum of the oxidation states = overall charge on ion

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8
Q

oxidation, reduction and oxidation states

A

oxidation = increase in oxidation state
reduction = decreased in oxidation state

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9
Q

disproportionation meaning

A

when the same element is oxidised and reduced in the same reaction

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10
Q

rule for combining half equations

A

make sure the number of electrons lost = the number of electrons gained

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11
Q

basic set up of electrochemical cells

A
  • two metal electrodes, each dipped in a solution of their own ions
    with a salt bridge connecting them
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12
Q

what is a salt bridge

A

a strip of filter paper soaked in a salt solution (e.g. KCl(aq))

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13
Q

setting up an electrochemical cell with two different metals process

A
  1. Take a strip of each metal and clean them with sandpaper
  2. clean them with propanone to remove oil/grease
  3. partially dip each strip into a beaker containing its ions
  4. add a salt bridge
  5. Use crocodile clips and wires to connect each electrode to the circuit or a voltmeter
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14
Q

how do electrochemical cells work

A
  • electrons flow from the most reactive metal to the least reactive metal, because more reactive metals will lose their electrons more easily (so oxidised more easily)
  • The salt bridge allows ions to transfer, completing the circuit
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15
Q

standard electrode potential definition

A

the potential difference between a metal and 1 moldm-3 solution of its ions, measured at 298K relative to the standard hydrogen electrode

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16
Q

what does it mean if the standard electrode potential is more negative

A

more reactive metal
provides electrons
stronger reducing agent

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17
Q

half cell with the most negative standard electrode potential

A

equilibrium shifts to left
half equation is reversed
becomes negative electrode
this half cell supplies electrons (reducing agent)
so oxidation occurs at this electrode, so this is the anode

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18
Q

half cell with the most positive standard electrode potential

A

equilibrium shifts to the right
half equation goes forward
becomes positive electrode
this half cell accepts electrons (oxidising agent)
so reduction happens here, so this is the cathode

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19
Q

calculating the cell potential from two standard electrode potentials

A

by subtracting the most negative one from the most positive one

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20
Q

predicting feasibility from electrochemical cells

A

First, determine which direction each reaction goes in
If reaction with more negative electrode potential is the reducing agent, then reaction is feasible

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21
Q

when would feasibility predictions be incorrect

A
  • when rate of reaction is too low
  • when activation enthalpy is too high
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22
Q

rusting definition

A

the corrosion of iron

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23
Q

conditions for rusting

A

oxygen
water
salt accelerates rusting

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24
Q

rusting overall equation

A

4 iron + 3 O2 + 2xwater –> 2 Fe2O3 .x H2O

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25
rusting mechansism
1. Fe(s) + 1/2 O2(g) + H2O(l) --> Fe 2+ (aq) + 2OH- (aq) 2. Fe 2+(aq) + 2 OH-(aq) --> Fe(OH)2(s) 3. Fe(OH)2(s) --oxygen + water--> Fe(OH)3(s) 4. Fe(OH)3(s) --oxygen + water--> Fe2O3 .xH2O(s) (hydrated iron (III) oxide)
26
prevention of rust
barrier method: paint, grease, oil, or less reactive metal galvanising: coating in a more reactive metal
27
transition metal criteria
- form complexes with ligands - form at least one ion with a partly filled d-subshell - d-block element - variable oxidation state
28
colours of transition metal ions in solution
Fe2+ - green Fe3+ - yellow/orange/brown Cu+ - unstable in solution Cu2+ - blue
29
ligand definition
a molecule or ion which can donate a pair of electrons (has lone pair(s) of electrons) to form a dative covalent bond
30
coordinate bonds definition
dative covalent bonds
31
complex ion meaning
a central metal surrounded by ligands, formed when transition metal compounds are present in a solution
32
central metal meaning
the metal the ligands bond to, can be an ion or atom
33
what is ligand substitution
the displacement of one ligand by another may change the coordination number of the metal, but not the oxidation state
34
relationship between electronegativity and ligands
less electronegativity makes for a better ligand as lone pairs are more easily donated so will displace weaker ligands
35
naming complex ions
1. prefix = number of ligands of the same type (tri, etc) 2. put ligands in alphabetical order if more than one type 3. name central metal 4. if overall charge of complex is negative, use latinised name of central metal (ferrate, cuprate, aurate) 5. indicate the oxidation state of central metal using roman numerals (II, III) 6. add the word ion on the end if it is one
36
monodentate ligand meaning
ligand that donates one pair of electrons (coordination number = number of ligands)
37
polydentate ligand meaning
when ligands donate more than one pair of electrons each (more than 1 coordinate bonds from the same ligand) (coordination number = number of ligands x number of coordinate bonds per ligand)
38
product when an aqueous transition metal ion and aqueous sodium hydroxide or aqueous ammonia are mixed
a coloured hydroxide precipitate
39
reaction and colours of Cu2+ (aq) + 2OH- (aq) (/dilute ammonia)
Cu2+(aq) + 2OH-(aq) --> Cu(OH)2(s) pale blue -------------------> blue
40
reaction and colours of Fe2+(aq) + 2OH- (aq)(/dilute ammonia)
Fe2+ (aq) + 2OH- (aq) --> Fe(OH)2(s) pale green --------> green
41
reaction and colours of Fe3+(aq) + 3OH-(aq)(/dilute ammonia)
Fe3+(aq) + 3OH-(aq) --> Fe(OH)3(s) yellow -------> brown / orange
42
reaction of excess NH3 + Cu2+(aq)
forms [Cu(NH3)4(H2O)2]2+ aq a dark blue ammonia complex
43
heterogeneous catalysts meaning
catalyst and reactants in different states
44
homogeneous catalysts meaning
catalyst and reactants in the same state
45
transition metals as catalysts
- transition metals have variable oxidation states allowing them to be good catalysts as it allows them to gain or lose electrons, providing an alternative route for lower activation enthalpy
46
transition metals as heterogenous catalysts
- Weak bonds are formed between the reactants and the 3d and 4s electrons on the catalyst - These bonds break after the product is formed - Transition metals are ideal because the bonds are strong enough to attract reactants, but weak enough for products to desorb
47
transition metals as homogeneous catalysts
- The catalyst usually forms an intermediate compound with one or more reactants - Later this will break down and the catalyst will be released again - Transition metal ions can move between different oxidation states easily - Because this enables them to be both reducing agents and oxidising agents, they are often good homogeneous catalysts - This is especially true for redox reactions
48
the colours of visible light
ROYGBV red orange yellow green blue violet
49
the amount of visible light absorbed by a molecule depends on
- the metal - its oxidation state - the ligand - the coordination number - the size of the energy gap
50
transition metal ions and colours
transition metal ions with a partially filled 3d subshell in solution will be coloured
51
what colour?
if an ion absorbs a particular colour, the observed colour will be the opposite on the colour wheel - the complementary colour
52
what do colorimeters measure
the absorbance of a solution
53
relationship between absorbance and concentration of a solution
Absorbance of a solution increases with concentration
54
what is a cuvette
a clear container which doesn't absorb light, used to hold the samples
55
method for using a colorimeter
1. Pick a filter with complementary colour to the colour of the solution 2) Put a sample of the solvent into the colorimeter and set it to zero 3) Put a sample in a clean cuvette. The absorbance is measured relative to the blank sample from step 2
56
Finding the concentration of a transition metal solution
complete colorimetry and measure absorbance of many samples with a known concentration and plot these on a calibration graph now can measure the absorbance of an unknown concentration solution and find the concentration from the graph
57
what is visual spectrometry
passing a beam of light with a specified wavelength through a sample, measuring how much is absorbed
58
key half equations in rusting process
Fe(s) --> Fe2+ (aq) + 2e- 1/2 O2 (g) + H2O(l) + 2e- -> 2OH- (aq)
59
what conditions is rusting less likely to occur in
alkaline conditions
60
61
why is rusting less likely to occur in alkaline conditions
half equations in rusting: 1. Fe2+ + 2e- <=> Fe(s) 2. O2 + 2H2O + 4e- <=> 4OH- in alkaline conditions, we have an increase in OH- ions, so the equilibrium in reaction 2 will shift to the left to remove extra OH- as a result there will be more electrons available because of this, equilibrium in 1 will shift to the right to use up extra electrons as a result, less Fe2+ ions, so less likely for rust to be produced
62
colour of agents in manganate titrations
MnO4 - is deep purple and is the oxidising agent Fe2+ is colourless (very pale yellow) and is the reducing agent
63
purpose of manganate titrations
to find out the concentration of Fe2+ or MNO4 - by titrating one against the other
64
what is the end point in manganate titrations
when there is a permanent colour change (either Fe2+ solution goes from colourless to pink, or MnO4- solution goes from purple to colourless)
65
what do we need to add to Fe2+ solution in manganate titrations and why
need to add excess dilute H2SO4 to ensure you have sufficient H+ ions to allow reduction of oxidising agent
66
what does electrode potential tell us about a half-cell
how easily the half cell gives up electrons (is oxidised) more negative half cell will undergo oxidation more positive half cell will undergo reduction
67
what is a ligand
an ion, atom, or molecule that ha at least one lone pair of electrons
68
monodentate ligand examples
H2O: :NH3 :CN :Cl
69
bidentate ligand examples
ethanedioate (lone pairs on Os single bondedly attached to C)
70
what is an electrochemical cell made up of
3 half cells joined by a wire, voltmeter and a salt bridge
71
flow of electrons in an electrochemical cell
flow from more reactive metal to less reactive metal

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