5.2.3 Redox and Electrode Potentials Flashcards
what is reduction
gain of electrons
- decrease in oxidation number
what is oxidation
loss of electrons
- increase in oxidation number
what occurs in redox reactions
oxidation and reduction
what is key when assigning oxidation numbers
- group 1,2,3 are always just their charge
- fluorine is always -1
- H is always +1, except for metal hydrides
- O is always -2, except with peroxides
- for ions, all the oxidation numbers always equal the overall charge number
what is an oxidising agent
takes electrons from the species being oxidised
- is REDUCED itself
what is a reducing agent
adds electrons to the species being reduced
- is OXIDISED itself
how can you form redox equations from half equations
1) balance out the number of electrons
2) add the equations together, and cancel out the electrons
3) cancel out all other species present on both sides
what is a way to balance out redox equations that involve many O’s and H’s
- via balancing changes in oxidation numbers
- rather than just trial and error
how do you form redox equations using oxidation numbers
1) assign oxidation numbers to all species
2) identify the species that experience a change in oxidation number
3) balance out ONLY the species that have changed in oxidation number(e.g. if on increases by 6, and one decreases by 1, multiply the decreased one by 6 to match the overall change)
4) balance out all of the remaining atoms
what may be missing in redox equations
you may not know all of the species involved in the reaction, so may have to predict products/reactants
what is a common reactant/product of aqueous redox reactions
- H2O is often formed
- H+ and OH- are also likely
- WHEN ADDINGS, ALWAYS DO A FINAL CHECK TO MAKE SURE BOTH SIDES OF THE EQUATION BALANCE BY CHARGE
how would you balance redox equations, where you are predicting products/reactants
1) assign oxidation numbers, and identify the species that change in oxidation numbers
2) balance out the electrons, by finding which side is more positive, so needs them added on
3) balance the remaining species, and add on your predicted products by using which elements are left out of overall equation
4) do a final balance of everything
ADD IN CARDS ABOUT REDOX TITRATIONS
what is a voltaic cell
a type of electrochemical cell which converts chemical energy into electrical energy
- takes place in modern cells and batteries
how do electrochemical cells generate electricity, and what does this mean for the type of reaction
- the electrical energy results from movement of electrons
- so requires a chemical reaction which transfers electrons from one species to another
- = redox reactions
what do half cells contain
the chemical species present in the redox half equations
how is a voltaic cell made
- by connecting 2 half cells
what is important about the 2 half cells in voltaic cells
- the chemicals in each one must be kept apart
- if they mixed, the electrons would flow in an uncontrolled way
- so heat energy would be released
- rather than electrical
how would you draw the Zn2+(aq)/Zn(s) half cell
- beaker containing solution of Zn2+
- a rod in the middle of Zn(s)
what are the two types of half cells
metal/metal ion
metal ion/metal ion
how is a metal/metal ion half cell set up
a metal rod dipped in a solution of its aqueous ion
- SIMPLEST, where both same element
what is the phase boundary of a metal/metal ion half cell
where the metal is in contact with its ion
- an equilibrium is set up here
what is the conventional way of writing an equilibrium
- FORWARD REACTION: shows reduction
- BACKWARD REACTION: shows oxidation
- e.g. Zn2+ + 2e- ⇌ Zn
what happens in an isolated half cell
there is no net transfer of ions either into or out of the metal
what is the direction of electron transfer dependent on when 2 half cells are connected
e- flow depending on the relative tendency at each electrode to release electrons
how are ion/ion half cells set up
- contain ions of the same element in different oxidation states as the solution
- there is no metal to transfer electrons into or out of the half cell, so an inert metal electrode is used, e.g. Pt
- shown via beaker with solution, e.g. Fe2+/Fe3+ and the electrode labelled with Pt
how do you find out which electrode has the greatest tendency to lose or gain electrons
- electrode with MORE REACTIVE metal: loses electrons and is oxidised, forming the NEGATIVE electrode
- electrode with LESS REACTIVE metal: gains electrons and is reduced, forming the POSITIVE electrode
what is the sign for standard electrode potential
E°
what is E°
the tendency to gain electrons, and be reduced
what is the standard chosen for E° values
- hydrogen gas
- and a solution of H+ ions
- with an inert platinum electrode
- the standard electrode potential of this half cell is 0V
what are the standard conditions needed for E°
- all solutions at 1 moldm-3 (e.g. the acidic H+ solution)
- 298K (25C) for gases
- pressure of 100kPa
how would you draw the standard hydrogen half cell set up
- beaker containing you H+ solution
- the platinum electrode
- a glass tube to allow H2 bubbles to escape
- the H2 gas entering from this glass tube at the side
- REMEMBER TO LABEL ON STANDARD CONDITIONS
what is standard electrode potential
the emf (voltage) of a half cell connected to a standard hydrogen half-cell under standard conditions
how do you measure the standard electrode potential
connect the half cell to a standard electrode (hydrogen)
how would you set up apparatus to find the E° of a half cell
- set up both of your electrodes
- connect them with a wire attached to the electrodes (to allow a controlled flow of electrons)
- have a voltmeter in the middle of this wire
- ALWAYS SHOW THE FLOW OF ELECTRONS IN THIS WIRE USING AN ARROW
- connect the solutions with a salt bridge (to allow the ions to move and flow)
what are typical examples of salt bridges
- usually contain a concentrated solution of an electrolyte that does not react with either solution on filter paper
- e.g. filter paper soaked in KNO3(aq)
where can you find the E° values of many redox systems
many have already been measured and listed in data reference tables
- REMEMBER: in data tables, they always show the equilibrium so that the forward reaction is reduction (ALWAYS GAINING ELECTRONS)
what does a more negative E° value show
- a greater tendency to lose electrons
- and undergo oxidation
- (the less the tendency to gain electrons and undergo reduction)
what does a more positive E° value show
- the greater the tendency to gain electrons
- and undergo reduction
what E° values do metals tend to have
- more negative E°
- so more likely to lose electrons
- the more negative this value is for a metal, the greater reactivity it has, in losing electrons
what E° values do non-metals tend to have
- more positive E° values
- so more likely to gain electrons
- the more positive this value is for non-metals, the greater their reactivity in gaining electrons
what is a cell potential
- the e.m.f value measured when connecting 2 half cells
- via symbol Ecell
- so just he potential difference between 2 half cells
how do you set up half cells to measure cell potentials
1) prepare 2 half cells at standard conditions
2) connect the metal electrodes of the half cells using a wire to a voltmeter
3) prepare a salt bridge (soak filter paper in saturated KNO3(aq)
4) connect the 2 solutions using a salt bridge
5) record the standard cell potential using a voltmeter
what is the flow of electrons in a cell potential Ecell
from the MORE negative half cell to the LESS negative half cell, so from the negative to positive electrode
- less negative E° is more likely to gain electrons
how do you write an overall equation from electrode potentials
- combine the relevant E° values:
1) REDUCTION equation stays the same, as if gaining electrons
2) OXIDATION half equation is reversed, to show the electrons are being lost
3) overall, combine the 2 half equations, and get rid of the electrons
how do you calculate standard cell potential
- using the standard E° values, and the difference between
E°cell = E°(positive electrode) - E°(negative electrode)
