Unit 4 Flashcards
Redox Reactions
reduction- oxidation
- when on or more electrons are transferred from one species to the other
oxidation
loss of electrons, oxidation number increases
reduction
a gain of electrons, oxidation number decreases
half-reactions
they are like writing one element and the electrons it looses/gains
e- is product
oxidation
e- is reactant
reduction
oxidizing agent
the one that gets reduced
reducing agent
gets oxidized
oxidation number
elements start out with an oxidation number of zero
for ions composed of only one atom, the oxidation number is the charge on the ion
oxidation number of oxygen
is mostly -2 even if its covalent
Exception:
Perioxides (R2O2) has a charge of -1
oxidation number of hydrogen
is +1 whether its ionic or covalent
exceptions:
when Hydrogen is bound to a binary metal ex. NaH, it has an oxidation number of -1
fluorine oxidation number
has -1, as do most halogens
exceptions to oxidation numbers
halogens other than fluorine is combined with oxygen to form an oxygen, the oxidation number becomes POSITIVE.
ClO- becomes +1 oxidation for Cl
sum of oxidation number
must give charge of compound
ex. ClO2-
Cl is +3
O2 is -4
electrons are -ve so…
if oxidation number for a specific atom becomes more +
- it lost electrons and is oxidized
if oxidation number for an atom becomes more -ve
- it gained electrons and is reduced
how to balance redox eqns
- write out half reactions
how to do half reaction
1) write unbalanced reaction in ionic form
2) separate ion into two reactions (one compound and its ions, and the other)
3) balance atoms other than o and H in each reaction separately, then add water to balance the O and H+
4) add electrons to one side of each half reaction to balance the charges
5) if needed, multiply both balanced reactions by an integer to equalize the number of electrons transferred in reactions
6) add them together and cancel identical species
converting chemical energy into electrical energy
when reagents are mixed, the electron transfer occurs directly and no electrical work can be obtained
electrical work
can be obtained through a movement of electrons through a wire.
- a spontaneous redox reaction produces electrical energy in a cell
energy cells
are galvanic cells
galvanic cells
1) solutions (cu and zn 2+?) are kept in separate containers called half cells.
2) electrode is dipped in each half cell. electrodes provide a surface for each of the chemical half-reactions to take place. these electrodes are made of the two metals involving Solid zinc, in zinc solution and solid copper in copper solution) or from a chemically inert metal such as plutonium
3) the oxidation container and redox container are connected by a wire. and the transfer of electrons occurs indirectly through the external wire
galvanic cell reaction notes
as reaction continues on, the solid zinc will get smaller (oxidation, anode, +)
as reaction continues on, the solid copper will get larger (reduction, cathode, -)
voltmeter is attached to wire to measure stuff
galvanic cell reaction without salt bridge (??)
as reactions take place, the zinc oxidizes and their are more ions in solution. the copper gets reduced and copper ions go into the metal. which would oppose electron flow and stop reaction as the anode is getting positive and the cathode is getting -ve
galvanic cell design must contain
species in solution and concentration, electrode and what it is made of, direction of electrons, voltmeter, salt bridge, labels of cathode/anote and reduction/oxidation and half reaction enqs
cell potential
- is cell electrical energy,
- measure of the force behind redox run
- E (cell potential) describes the ability of each reagent to donate or withdraw electrons
cell potential depends on
chemical reaction and concentration of solutes and/or partial pressures of gases
voltmeter
- measure cell potential
- requires the flow of electrons which is slightly lower
- part of the cell potential is used to overcome the resistance of the wires in the voltmeter
- will measure a potential that is less than the maximum cell potential
shorthand notation
phase boundaries are represented by a: /
reactants are written before products
reduction is written on the right
ex.
Zn (s)/ ZN2+ (aq, 1 M) // Cu 2= ..
Reactants/Products//Reactants/Products
the double slash is the salt bridge.
oxidation number notation
write the charge BEFORE THE number
ex. -1 vs 1- (this is charge)
halogen-Ox ion charge (ex. ClO3-)
this will always have a charge of 1- overall
standard electrode potential, E cell
reduction potential for each half cell can be found.
E cell
E cell= E red + E ox
this describes the ability to gain or lose electrons compared to the standard hydrogen electrode
standard hydrogen electrode (SHE)
you can’t measure the reduction potential of a single half cell since one half cell has to be connected to another half cell to work
solution was a relative scale by choosing the reduction potential of H+ as 0.000V
2H+ + 2e- –>—< (equilibrium) H2
the She device
electrode is made of platinum, solution is of H+ and H2 gas is made so an inverted beaker or something captures the gas.
H+ is measured as pH to see whether a reduction or oxidation took place as high pH would mean low H+, equilibrium would shift right which is reduction.
reduction of H+ is reversible so the other half cell determines whether h+ is reduced or H2 oxidized. the evidence for this is the increase in H2 gas, high pressure or effervescence.
Calculating cell potential
half cell with larger E red (greater potential for reduction) will be reduced
the other will be oxidized
e ox= -e red
salt bridge
without the salt bridge:
solution is neutral at the start but the anode gets more +ve charge and the cathode is getting more -ve charge which is opposing electron flow.
With bridge:
the ions like K+ and NO3- seep into solutions and keep them neutral.
anode charge
is negative
cathode charge
is positive
electron flow is from
anode to cathode
go draw a galvanic cell with zinc and copper
ah
