Redox II Flashcards
Oxidation in terms of electrons and oxidation number
Loss of electrons
Increase in oxidation number
Reduction in terms of electrons and oxidation number
Gain of electrons
Decrease in oxidation number
Always write redox equilibria int he form of
Reduction (electrons on LHS on equation)
What happens to electrons when a metal is placed in water?
Metal atoms shed electrons and the metal ions go into the water so that the electrons build up on the surface and result in a negative charge
How are positive ions attracted to a metal in water?
Buildup of electrons on the metal attracts positive ions in the solution to form a layer of positive ions
How do you some of the positive ions in solution form part of the metal?
Some of the positive ions in the layer regain that electrons from the buildup of electrons and become part of the metal surface again as a metal atom
How is a dynamic equilibrium establish when a metal is placed in water?
When the rate at which ions are leaving the surface = rate at which they are joining again (metal atoms reforming)
Explain what the diagrams for Magnesium in water and Copper in water as a snapshot of dynamic eqm look like?
Magnesium: lots of electrons on the surface and big layer of positive ions
Copper: less electrons on surface and less of a layer of positive ions (because Copper is less reaction so the Cu atoms shed electrons to form positive ions less easily)
Position of equilibrium in relation to how well something sheds slew from a and forms positive ions?
Readily sheds electrons and forms positive ions
Equilibrium lies further to the left
Standard Hydrogen electrode
Equilibrium Pressure Temperature Concentration Catalyst
2H+(aq) + 2e- H2(g) Hydrogen pressure at 1 bar (100kPa) 298K H2SO4 conc is 1mol dm^-3 Platinum catalyst
What is a cell?
The whole set up (two electrodes connected by a voltmeter and saltbridge)
Aka two half cells
What is a half cell?
Each of the two beakers and their contents
Ie. Mg electrode in beaker of Magnesium Sulphate
What is the purpose of a salt bridge?
Included to complete the electrical circuit without introducing any other pieces of metal into the system
What is a salt bridge?
Glass tube filled with an electrolyte like potassium nitrate solution
End stoppered by cotton wool to prevent mixing of contents in salt bridge with those in beaker
What will the system look like if a Magnesium electrode and platinum electrode for hydrogen are measured?
Magnesium electrode had much greater build up of electrons that the platinum electrode
Mg eqm lies further left
Explain why the voltmeter must be high resistance?
Avoids flow of current through the surface
What does E cell mean?
The electromotive (emf) of the cell is the maximum possible voltage in any situation.
What is the cell diagram for when a magnesium electrode is coupled to a hydrogen electrode
And variants
Pt [H2(g)] | 2H+(aq) || Mg2+(aq) | Mg(s)
Pt | H2(g) | 2H+(aq) ||
Or
Pt | H2(g) | H+(aq) ||
X
X
Cell diagram conventions (3)
Draw an arrow from the E standard to the right hand electrode
[] Shows something flowing over a catalyst (eg. Hydrogen gas over platinum)
Substance losing electrons is written closest to its electrode
If more than one thing on either side of equilibrium (eg. With potassium dichromate(VI)) square brackets written around to keep tidy
Define standard electrode potential
Emf measured when a metal electrode (or a metal ion electrode) is coupled to the standard hydrogen electrode under standard conditions
Explain why equilibrium lying furthest to the left means the most negative E value?
Form ions more easily leaving electrons behind on the metal
making it more negative
Connect E values, equilibrium direction and readily forming ions
Negative E value
Equilibrium further left
More readily element loses electrons to form ions
Stronger reducing agent
Positive E value Equilibrium further right Less readily loses electrons/forms ions/readily picks up electrons again Stronger oxidising agent
What is the electrochemical series?
Arranging various redox equilibria in order of standard electrode potentials
Most negative E value at top
Most positive E value at bottom
Explain strong oxidising and reducing agents in relation to the electrochemical series?
Top: form ions more easily, good at giving away electrons so GOOD REDUCING AGENTS
Bottom: form ions less easily, good at picking up electrons so GOOD OXIDISING AGENTS
Oxidising and reducing ability in relation to electrochemical series?
Oxidising ability of ions increases down the electrochemical series
Reducing ability of the metal increases up the electrochemical series
Equilibrium lying more to the left means…
more likely to pick up electrons
What happens if you have a system involving only gases?
Eg. Hydrogen and chlorine
Set up standard hydrogen electrode as normal
Set up chlorine gas bubbling over another platinum electrode immersed in 1 mol dm^-3 Cl- solution
Explain how to measure redox potential for a system of Fe2+/Fe3+ ions
Platinum electrode inserted into beaker containing solution of iron (II) and iron (III) ions coupled with hydrogen electrode
Potassium dichromate (VI) equilibrium
Cr2O72- (aq) + 14H+ (aq) + 6e- 2Cr3+ (aq) + 7H2O(l)
Potassium dichromate half cell set up
Standard hydrogen electrode coupled with
Platinum electrode in beaker with solution containing Cr2O7 2-(aq), H+(aq) and Cr 3+(aq) ions at 1 mol dm^-3
More negative E value means
Release electrons and forms ions more readily
Zinc half cell and a copper half cell
You can combine two half cells not including the standard hydrogen electrode
The more negative E value = forms ions more readily = eqm further left compared to other half cell
Copper: less electrons and layer of positive ions
Zinc: more electrons and larger layer of positive ions
Explain how to show the relationship between combinations of half cells and simple redox reactions carried out in test tubes
1) write redox equilibriums of both electrodes and E values
2) work out which E is the most -ve value so can work out which direction electrons flow
3) work out which way this means the eqm’s shift
4) write two one way redox equations
* see notes*
Examples of test tube reactions that can be shown by redox equilibria electrodes without voltmeter (chemguide) (4)
1) Combining zinc half cell and copper half cell (e- flow from zinc to copper)
2) reaction between copper and silver nitrate (e- flow from copper to silver)
3) Reaction between magnesium and dilute sulphuric acid (e- flow from magnesium eqm to hydrogen eqm)
4) potassium dichromate (VI) oxidising Fe2+ —> Fe3+ (e- flow from Iron eqm to dichromate eqm)
In whole cells where the voltmeter is removed
what is the rule of thumb for which way eqm shifts with E value?
Equilibrium with more negative E value shifts to left
Equilibrium with more positive E value shifts to the right
How do you use an E value to predict the feasibility of a possible redox reaction?
1) write out the two redox equilibria and E values
2) find the most -ve E value to work out which direction the electrons flow. Then work out what effect this has on the shifting equilibrium.
3) check what things you start with and if the shift of equilibrium can happen versus if it can’t shift any more in one direction to work out feasibility/non feasibility
Which redox reaction is predicting the feasibility difficult for depending on what reacts?
Copper and dilute nitric acid
Copper could react with either H+ (not feasible because eqm already shifted as far as possible)
or
NO3 - (feasible) and produces Cu2+ and NO2
Why might an E value predicting feasibility appear incorrect?
X
Not feasible because large activation energy barrier
Name two examples where E values predicting feasibility appear wrong and say why.
1) acidified potassium dichromate (VI) oxidising water.
E value says it is feasible but nothing happens in the test tube because a large activation energy barrier must be overcome
2) acidified potassium dichromate (VI) oxidising chloride ions to chlorine
E value says it is not feasible however it is feasible if you change standard conditions using HCL to give H+ and Cl- ions at 10 mol dm^-3 concentration
Explain how to use redox equilibria to select an oxidising agent
Write out redox equilibria, in reduction form, and E value for thing you want to oxidise
Eg. Fe3+(aq) + e- Fe2+(aq). E = +0.77v
To oxidise equilibrium must shift left (E value for redox equilibrium you want to shift must be more negative)
So find another equilibrium you can couple this with that has an E value that is less negative than the eqm one you already have
Explain how to use redox equilibria to select a reducing agent
Write out redox equilibria, in reduction form, and E value for thing you want to reduce
Eg. Cr3+(aq) + e- Cr2+(aq). E = -0.41
To reduce, equilibrium must shift right (E value for redox equilibrium you want to shift must be less negative)
So find another equilibrium you can couple this with that has an E value that is more negative than the eqm one you already have
