redox Flashcards
describe reduction in terms of electrons
gain of electrons
describe reduction in terms of oxidation number
decrease in oxidation number
describe oxidation in terms of electrons
loss of electrons
describe oxidation in terms of oxidation number
increase in oxidation number
state the oxidation number for an element
0
state the oxidation number for combined H
+ 1
state the oxidation number for combined O
- 2
state the oxidation number for an ion
equal to the ion’s charge
describe an oxidising agent
takes electrons from another species (oxidises it)
is reduced
describe a reducing agent
adds electrons to another species (reduces it)
is oxidised
state the oxidising and reducing agent in the following equation:
2Ag+ + Cu –> Cu2+ + 2Ag
Ag = oxidising agent
Cu = reducing agent
how are redox equations balanced?
by balancing the electrons across both equations
balance the following redox equations
H2O2 + 2e- –> 2OH-
Cr3+ + 8OH- –> CrO4 2- + 4H2O + 3e-
3H2O2 + 6e- –> 6OH-
2Cr3+ + 16OH –> 2CrO4 2- + 8H2O + 6e-
describe the process for balancing half equations
write species before and after reaction
balance atoms except O and H
balance O with H2O and H with H+
balance charges with e-
balance the following equation
Fe2+ –> Fe3+
Fe2+ –> Fe3+ + e-
balance the following equation
MnO4- –> Mn2+
MnO4− ( a q ) + 8H+ ( a q ) + 5e− → Mn2+ ( a q ) + 4H2O ( l )
in a redox half equation, where are the electrons positioned for reduction
on the right
in a redox half equation, where are the electrons positioned for oxidation
on the left
describe the process for combining half equations into a full ionic equation
balance electrons
cancel electrons
combine
state the oxidation number for oxygen in fluorides
+ 2
state the oxidation number for oxygen in peroxides
- 1
state the oxidation number for hydrogen in metal hydrides
- 1
define oxidation number
the charge that a molecule would have, if the electrons in each bond belonged to the more electronegative element
