deck Flashcards
current definition of oxidation
complete or partial LOSS of electrons (and gain of oxygen)
current definition of reduction
complete or partial GAIN of electrons ( and loss of oxygen)
which side of the equation are the electrons in an oxidation equation?
products (losing electrons)
what happens to amounts of oxygen, hydrogen, and electrons in a reduction reaction
lose oxygen (REDUCTION in the amount of oxygen)
gain hydrogen and electrons (they balance each other out)
what is oxidized, reduced, and spectating in this equation
2AgNO3(aq) + Cu(s) –> Cu(NO3)2(aq) + 2Ag(s)
oxidized - copper, specifically Cu(s) (RA)
reduced - silver, specifically Ag+(aq) (OA)
spectator - NO3-(aq) (dissociates into the same ion form on both sides)
when and why are oxidation numbers used
used to identify if covalently bonded elements gain or lose e- because they share electrons so it is only a PARTIAL gain/loss
what is the oxidation number of pure monatomic and polyatomic elements
what is it for H2
zero for both
oxid # for H2 is 0
what is the oxidation number of a monatomic ion
what is it for Mg2+
equal to its charge
oxid # for Mg2+ is +2
what is the oxidation number of group 1s, group 2s, and (most) group 17s
what is it for F
(equal to their charges on the periodic table)
group 1 is +1
group 2 is +2
group 17 is -1
oxid # for F is -1
what is the oxidation number of H and what is the exception
usually +1
exception is ionic hydrides where H acts as the cation. here the oxid # for H is -1
ex. NaH
what is the oxidation number of O and what is the exception
usually -2
exception is peroxides, where there are 2 oxygens where it would typically make more sense to have one. here the oxid # for O is -1
ex. H2O2
notice that the charge is still neutral, so if the oxid # for O was -2 like usual, the charge wouldn’t balance because the oxid # for H is +1
what are the oxidation numbers in each element in K2SO4
K is +1
S is +6
O is -2
O is always -2, which is multiplied by the subscript. K is in group 1 so its charge is +1 which is multiplied by the subscript. with those two elements, the charge of the compound is -6, but has to be neutral so S is +6
what do oxidation numbers tell us
when it increases from reactants to products, oxidation occurs, making that element a reducing agent
when it decreases from reactants to products, reduction occurs, making that element an oxidation agent
what are the oxidizing and reducing agents in this equation
Cl2(g) + 2HBr(aq) –> 2 HCl(aq) + Br(l)
reducing agent (is oxidized) = Br-
oxid # -1 –> 0, increases
oxidizing agent (is reduced) = Cl2
oxid # 0 –> -1, decreases/REDUCES
general rules for sample and titrant in redox reaction titrations (consider whether its oxidized or reduced, if the concentration is given, and which specific compounds or ions can be used)
SAMPLE
- oxidized (RA)
- unknown concentration (you will be asked to solve for this)
- usually the lower valence charge of a multivalent metal
TITRANT
- reduced (STRONG OA)
- known concenttration
- have to CHANGE COLOUR when they reduce so we can see when to stop the titration
steps (with formulas) for a titration calculation (concentration of sample)
- measured volume of sample and initial and final volumes of titrant (average)
- calculate moles of titrant used using n=cV
- calculate moles of sample using the molar ratio
- calculate concentration of sample using c=n/V
how to balance redox rxns
- break into half reactions, do steps 2-5 on each reaction
- balance all except H and O
- balance oxygen first by adding enough H2O to the side with less to balance on other side
- balance hydrogen by adding H+ ions to the side with less (include H2Os added last step)
- add electrons to the side with a more positive charge (usually opposite side you added H+ to)
- multiply by LCM so that each equation has an equal number of electrons, then recombine
- double check by making sure everything (atoms, number of atoms, charges) is balanced
electrons lost in oxidation
gained in reduction
and transferred in total in this equation:
4NH3(g) + 7O2(g) –> 4NO2(g) + 6H2O(g)
oxidation - lost 7e-
reduction - gained 4e-
total - transferred 28e-
how to use the shortcut method to find how many electrons lost/gained in oxidation/reduction in a redox rxn
(change in oxidation number) x (subscript of reactant) / (subscript of product)
usually there is no subscript for the product
how to use the shortcut method to find how many electrons transferred in total in a redox rxn
(not sure if this is correct but it should work at least most of the time??)
the number of electrons transferred in total is the lowest common multiple between electrons lost in oxidation and gained in reduction
what is disproportionation and how do we recognize it
when a species is both oxidized and reduced
recognize when the same species is found in two places (2 different cmpds or ion) on the products side
ex. H2O2(l) –> H2O(g) + O2(g)
oxygen disproportionates. with one product the oxid # increases, with the other it decreases, meaning it both oxodized and reduced
true or false, a battery is a non spontaneous reaction because they don’t react until connected to an appliance
false, they are spontaneous. the reason they don’t react is that the OA and the RA are separated until e- can flow through the appliance from the RA (a metal) to the OA
what energy conversion happens in voltaic cells
chemical Ep –> electrical energy
4 things you need in a voltaic cell
- electrode (2) - carries electrons to or from a substance
- electrolyte (2) - conducting solution that contains the same ions as the metal in its electrode, needs cations and anions throughout
- wire - allows e- movement from the anode to the cathode
- porous boundary - either a salt bridge or a porous cup, allow flow of charged ions to balance overall cell charge
how do you know which electrode in a voltaic cell is the anode vs cathode and which way do electrons flow
oxidation occurs at the anode (both vowels) (negative electrode)
reduction occurs at the cathode (consanants) (positive electrode)
since oxidation is losing electrons and reduction is gaining, the electrons flow from the anode to the cathode (alphabetically)
describe ion flow in a voltaic cell and the impact it has on electrodes physical properties
anions migrate towards the anode
cathodes migrate towards the cathode
since the anodes and cathodes generally submerged in their ion in aqueous form and that ion is a metal so its positive, those ions tend to flow through the boundary to the cathode. the cations flow this way to balance the charge of the whole cell because the electrons are also flowing to the cathode, so to keep it balanced the metal anion dissociates over time into its ions and the opposite happens at the cathode simultaneously to balance the charge, resulting in LOSS OF MASS at the ANION and GAIN OF MASS at the CATION
when can a salt bridge not be a nitrate
in acidic cells because NO3-(aq) + H+(aq) is a SOA (found at .80V on table)
when can a salt bride not be a chloride
in cells with silver or lead ions because it will form a precipitate
when can a salt bridge not be a sulfate
in cells with silver or lead ions because it will form a precipitate
formula for E°net
E°net = E°rpc - E°rpa
(1 value - another value, no multiplying)
reduction potential at the CATHODE - reduction potential at the ANODE = electric potential difference
qualitative observations in a lab for an electrolytic cell you should notice
- mass gain at cathode (if its a metal)
- loss of mass at anode (if its a metal)
- pH change (increase of H+ in products, decrease if in reactants)
- colour change of aqueous solution (production or consumption of a coloured ion, refer to page 11 data book)
define electrochemical cells
just the umbrella term for voltaic and electrolytic cells
define electrolytic cells
a non spontaneous reaction converting electrical energy to chemical Ep (reverse process of a voltaic cell
require a current to be added
chemical change THROUGH the application of electrical energy
used to purify metals
name the electrochemical cell for each idea
spontaneous
exothermic
like charging a phone
voltaic
voltaic
electrolytic
name the electrochemical cell for each idea
requires a power source
is a power source
negative Enet
electrolytic
voltaic
electrolytic
name the electrochemical cell for each idea
SOA reacts with SRA
connected to a voltmeter
connected to a battery
both
voltaic
electrolyitc
name the electrochemical cell for each idea
turns electrical energy into chemical
single container (no boundary)
operates based on both oxidation and reduction
electrolytic
electrolytic
both
describe electrolysis of water
consider:
- what is chemically being formed from water
- what is produced at anode and cathode
- acidity at anode and cathode
- what can the SOAs and SRAs be
- electrolyte
- E°net
splitting water into H2(g) and O2(g), which are visible as bubbles
(red. at cathode) the production of H2 also produces OH- so this area is basic
(oxid. at anode) the production of O2 also produces H+ so this area is acidic
water MUST be both the SOA and the SRA, meaning it will not happen correctly if there is a stronger OA or RA than water (there aren’t many reactions on the table weaker than water so be careful…)
added electrolyte to allow electron flow
E°net = E°rc - E°ra
= (-.80V) - (+1.23V) = -2.06V
(reduction of water) - (oxidation of water)
describe electrolysis of molten sodium chloride
consider:
- states!!!
- what is formed at cathode/anode
- purpose
- why electrolytic
- is it the same E°net with different formulas?
- E°net
(oxid. at anode) chloride ions form Cl2(g)
(red. at cathode) sodium ions form Na(l)
a major source of chlorine gas
needs to be an electrolytic cell because otherwise once separated the Cl2(g) and the Na(l) would form NaCl again
same E° applies, but when writing formulas we must write the altered states, and consider that there is no water present (can’t be aqueous anyway)
E°net = E°rc - E°ra
= (-2.71V) - (+1.36V) = -4.07V
(reduction of sodium) - (oxidation of chlorine)
describe the electrolysis of brine (saltwater)
consider:
- products
- purpose of creating these products
- what is produced at anode and why
- what is produced at cathode and what else is produced as a result
very common, produces Cl2(g) for plastics/agriculture, H2(g) for fuel, and NaOH for industrial cleaners. produced simultaneously
(oxid. at anode) chlorine gas produced at anode, THIS IS AN EXCEPTION, water “should” be oxidized here. (the Ea for chlorine in this one case is less than the Ea for water, you don’t need to know that though)
(red. at cathode) water is reduced to form hydrogen gas, also produces OH- which reacts with sodium ions to produce NaOH. sodium is not reduced here because water is SOA
describe electroplating
consider:
- what happens chemically
- what is the cathode and what happens there
- what is the anode
- what is the electrolyte
- purpose
- a metal ion is reduced to its solid metal form
- the cathode is the object getting plated (because the metal cations flow to it and gain the e-, becoming the solid metal) and e-s are entering into the cathode, REDUCTION
- anode is usually what is doing the plating but can also be inert
- performed in an electrolyte (solution) of the metal ions that are to be plating the cathode
- protects from corrosion, makes things look nice
- plating with chromium is special and is more resistant to oxidation from environment than other metals
*note - even when a reaction is not shown in the data book, you can infer it by adding the same amount of e- as the charge of the ion
describe corrosion
consider:
- what it is + is caused by
- what does it produce
- when is it a problem
- and why
- the breakdown of metals via e- transfer which is caused by exposure to O2(g) (it’s like dissolving but metals)
- rust is Fe(OH)2(s) in chem 30 (but there are further steps)
- only a problem in iron because metal oxide is impermeable to O2(g) and H2O(l), the oxide layer usually stops oxidation because only the outermost layer oxidizes
- all iron will continue to rust in the presence of ANY iron oxide
describe the three ways to prevent corrosion
- providing a barrier via painting, oiling, dipping, electroplating, alloying
- sacrificial anode (providing alternative e- source) by adding another metal that is a SRA than Fe(s), however the alternative metal needs to be replaced every so often once it runs out of e-
- impressed current (providing alternative e- source) where a current is sent through the object to provide e- and the Fe(s) will never lose its e- to oxidation because the O2 / H2O take e- from the Fe but the current instantly replaces it
which law allows you to calculate the mass of electroplated metal onto the cathode/lost at the anode?
Faraday’s Law (allows to solve for moles but using n=m/M we can find mass)
what are the units for a coulomb (C)
A*s
as in amperes x seconds
what are the units for faradays constant
(A)(s)/n
ampere x second / mole
