4.3 Oxidation Reduction Reactions Flashcards
The loss of electrons
Oxidation
so you get more positive
Reduction
The gain of electrions
* makes since, if you gain electrons you lose charge
In element in its atom form is in its 0 oxidation phase. - meaning it hasnt lost or gained electrons
Were looking at the change in oxidation state as evidence that electrons are either being lost or gained.
Reduced = losing charge, but gaining electrons
Oxidized = gaining charge, but losing electrons
Whenever you have an oxidation reduction rxn you’re always going to have at least one oxidation and one reduction
* technically you can have more
Oxidation rules
1) Atoms in their elemental form are in the zero oxidation state
* This makes sense, they have no net charge. They have neither lost, nor gained electrons - this is true of diatomics as well –> they must have no charge on them - aka be in their elemental form
2) Monatomic ions in ionic compounds get their typical ionic charge
* Think NaCl - these are both monoatomic ions (not polyatomic) - Na+1, Cl-1 (Na is oxidized, Cl is reduced)
* NOTE: if you have transition metals they can take on a variety of charges and you’ll have to figure them out based on what anions they’re attached to.
3) Atoms in Molecular Compounds and Polyatomic IOns
* Oxygen -2 except in peroxide O2^2- (where it gets a 1- charge)
* Hydrogen +1 (when bonded with any nonmetals)
* Fluorine = -1 always
* Halogens = -1 except when bonded to oxygen
4) The oxidation numbers sum to zero for a compound or the overall charge of a polyatomic ion
NOTE any neutral molecule must stay neutral, so if theres just 1 atom left in the compound you can give it the charge that would make the rest of the molecule neutral
how to recignize a peroxide
* it has 2 oxygens and its with either hydrogen, or a group 1 metal
* A peroxide contains tha O2^2- ion, where two oxygen atoms are covalently bonded to each other and carry a -2 charge overall. Each oxygen in the peroxide ion has a -1 oxidation state (unlike in most compounds where oxygen is -2)
* It typically looks like K2O2 or H2O2
* The key feature: O-O single bond and each oxygen has a -1 oxidation state
* So it makes sense that it has to be a 1- charge on the oxygen if you break this compound apart
Which one of the below is a redox reaction
* hall mark is the transfer of electrons - someone gains electrons someone loses them
* you could start by just going through an assigning oxidation states to all the compounds - but this is time consuming
Tricks:
* Look for elements in their elemental form
* For the first equation below you can see Mg and O2 both in their elemental form –> and if they arent in their elemental form on the other side, thats going to mean they arent in their 0 oxidation state
For the last one Cu goes from elemental form to part of a compound = a change = redox rxn
* This is actually a single replacement rnx - and single replacement rnxs are always redox reactions
** The third one is a double replacement rxn, and double replacement rxns are enevr redox - thats because we balance the ionic compounds and the electrons don’t actually change**
Example of double replacement not changing oxidative state
naCl + AgNO3 –> NaNO3 + AgCl
* Na+, Cl-, Ag+, and NO3- keep the same oxidation states before and after
* because they completely balance out
Single replacement: A free elemental replaces another element in a compound
* This usually means one element is oxidized and the other is reduced
Zn + CuSO4 –> ZnSO4 + Cu
* Zn goes from 0 –> Zn^2+ –> Oxidized
* Cu^2+ becomes Cu0 –> reduced
Single replacement Reactions = type of redox rxn
* In single replacement the alone reactant switches w/ the coupled reactant in the products
First the first one think of it like this:
* H and Cl are a couple, and Mg wants to be paired w/ Cl. In the products you can see that happens and H is left all alone.
* However, the only way for this to happen is if Mg is better looking than H.
* We can decide whose “better looking by whose higher up on the activity series” - meaning the higher up on the activity series = oxidized more easily = spontaenous rxn
To be oxidized more easily means that the compound loses electrons more readily than another compound. In other words, its more willing to give up its electrons and undergo oxidation
This is the activity series
Most active (reactive) metals in an oxidation rxn at the top, and the least reactive at the bottom
* They get more easily oxidized as they go up the list
Li –> Li+ + e-
* got oxidized
* This reaction is more easily going to happen according to the list below. Li is more easily oxidized (meaning it gives up an electron) than K
K –> K+ + e-
* Got oxidized
Dont memorize this –> just going to need to know that if its higher up on the list, its more easily oxidized (meaning it will give up an electron and turn positive more easily)
0
Diamond is an elemental form of carbon and any element in an elemental form is in the zero oxidation state (meaning its neutral)
oxidation rules
Trying to balance this didnt make since, thats because oxygen is a peroxide here
ways to recignize a peroxide:
* it has 2 oxygens and its with either hydrogen, or a group 1 metal
This is an ionic compound and when its ionic they keep their original charges
1-
+4
So I was able to see that K+ was a monoatomic ion so it retained its original charge
Oxygen when in a polyatomic gets a -2 charge
By knowing those 2 I was able to solve for Sulfur
+3
followed molecular compound rules
tricky part was at the end where I had to realize it was 2 carbon atoms and divide by 2 to get the oxidative state of the single carbon atom
key was remembering that it wasnt a neutral compound
single replacement
single replacement
The reason something higher up on the activity series (is oxidized more easily meaning it gives up its electrons more easily) causes a spontaneous reaction is because the other atom wants electrons and takes them from the one whose offering them more easily
Higher on activity series = gives up electrons more easily (oxidized)
Lower on activity series = accepts electrons more easily (reduced)
