Section 10- Developing Metals Flashcards
(32 cards)
In electro chemical cells, where do the following ALWAYS occur at?
-Oxidation
-Reduction
This is different to the standard
-Oxidation ALWAYS occurs at the Anode (-ve)
-Reduction ALWAYS occurs at the cathode (+ve)
The cathode and anode are usually the opposite charges!
What can voltmeters do in electrochemical cells?
Measure the potential difference between the two half cells
It can also measure the direction of the flow of the electrons, which is telling of which metal is the cathode/anode
In electrochemical cells, there can be half cells of ions of the same element. Give an example of this and explain how they work, including relevant apparatus.
Fe3+/Fe2+
The ions will oxidise and reduce into each other on the surface of a conductive and inert electrode (i.e. platinum or graphite)
When preparing the metal samples which are under investigation, what must you do to clean the strips to prevent unwanted contamination?
Clean the surface with emery paper (sand paper)
Use propanone to remove any grease or oils.
Once these are done, avoid touching the metal and where you’ll potentially transfer grease onto the metal.
Although half cell reactions are reversible, they’re ALWAYS written in a specific way. In what way are they written?
With the reduction reaction going forwards.
Cu2+ + 2e- –> Cu
What is the rule with electrode potentials in terms of the flow of electrons?
The most positive electrode potential are the ones which are reduced/the reaction goes forward:
Cu2+ + 2e- —> Cu
The most negative therefore goes in the opposite direction and is oxidised to provide these electrons.
What is different about the standard hydrogen electrode?
It is always shown on the left, regardless of if the other electrode in oxidising.
I think this is just to set as a standard referencing point.
Equation to calculate the electrode potential of an entire cell?
E°cell = E°most positive - E°most negative
E°cell = E°cathode - E°anode
What are the features of a REACTIVE METALS standard electrode potential?
They are LARGE AND NEGATIVE, meaning that they have a strong tendency to be OXIDISED/LOSE ELECTRONS into their metal ion form.
What are the features of a REACTIVE NON METALS standard electrode potential?
They tend to have LARGE POSITIVE electrode potentials, meaning that they have a strong tendency to be REDUCED/GAIN ELECTRONS to form negative ions.
Order you should use when writing redox equations:
Write the half equations
If there’s an oxygen involved, add a water to balance the oxygen
Then balance the hydrogen with H+ ions on the opposite side
Then weigh up the charges and add electrons to ensure a balanced charge
Then change molar quantities to make it so electrons cancel out
You can now combine into a complete redox equation
During redox titrations, what must you add to the reducing agent ensure the oxidising agent (i.e. MnO4-) is reduced?
You must add H2SO4 in excess as this ensures that there’s enough H+ for the oxidising agent to be reduced (they are required to balance charges from previous flash card)
Standard Electrode potentials can be used to predict whether the reactions will occur. Sometimes the predictions can be wrong. In what cases may this be apparent?
When the rate of reaction is so slow that the reaction may not appear to occur.
When the activation enthalpy is simply too high, and the reaction will not occur.
When the conditions aren’t standard (when the temperature or concentrations are changed, resulting in a shifted equilibrium)
What are the steps which occur to form rust?
Fe —> Fe2+ + 2e-
2H2O + O2 +4e- —> 4OH-
Fe2+ + 2OH- —> Fe(OH)2
2H2O +O2 + 4Fe(OH)2 —> 4Fe(OH)3
This then eventually forms rust:
Fe2O3.xH2O
What conditions make rusting less likely to occur?
Why?
Alkali conditions
Alkali conditions increase the concentration of hydroxide ions, which affects the position of the equilibrium, meaning that more solid iron is produced to counteract the increase in electrons. Iron(II) ions are needed to produce iron(II) hydroxide which leads to rusting etc.
Ways to prevent rusting?
Oiling and greasing- this involves coating the iron with a layer in which water and oxygen cannot resonate to react, so rust doesn’t form.
Coating with a polymer- this is decorative and ideal for large structures and small structures.
Sacrificial protection- This involves using a metal which is more reactive than zinc, where it will react rather than the metal one wishes to preserve. This can be done by galvanising (spraying a coat of zinc) an object, or rather placing a block of zinc onto a the metal. This is useful for ships or underground pipes.
What are transition metals?
Elements/metals found in the d block.
Can form at least one ion with an incomplete/split d orbital.
They have VARIABLE oxidation states.
Form coloured compounds
Although in the d block, what are not transition metals?
Scandium (Sc).
Sc3+ has an empty d subshell
Zinc (Zn)
Zn2+ has a full d subshell
Copper exists with oxidation states +2/+1, why don’t you see the +1 often?
It’s unstable in solution and disproportionate.
2Cu+ —> Cu2+ + Cu
What colour does each ion form in solution?
Fe2+
Fe3+
Cu2+
Fe2+ forms pale green
Fe3+ forms yellow
Cu2+ forms blue
What is a complex ion?
A complex ion is a complex where a central transition metal ion is surrounded by ligands, where coordinate bonds form.
What is a ligand?
A ligand is an atom/ion/molecule that donates a pair of electrons to a central transition metal ion, where a coordinate bond will form.
Ligands have different names depending on the number of coordinate bonds they form with the central transition metal ion. Name these:
1 coordinate bond with one lone pair
2
2+
Monodentate
Bidentate
Polydentate
What shapes do the following form?
Coordinate number:
6
4
2
Octahedral
Tetrahedral/square planar
Linear
