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LC Chemistry > Oxidation Numbers > Flashcards

Oxidation Numbers Flashcards

1
Q

Oxidation number

A

The oxidation number of an element is the charge an element has or appears to hav when it is ina compound when certain rules are applied

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2
Q

rules of oxidation numbers

A

8 rules

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3
Q

rule 1

A

Simple elements i.e. those that are not with any other element have an oxidation number of zero

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4
Q

rule 2

A

In the combined state, group 1 elements always have an oxidation number of +! and group 2 elements lways have an oxidation number of +2

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5
Q

rule 3

A

in simple ions, i.e. single element ions, the oxidation number of the ion is equal to the charge on the ion

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6
Q

rule 4

A

the halogens (group 7) in compounds where the halogen is bonded to ONE other element, has an oxidation number of -1. unless the halogen is bonded to a more electronegative element, then the oxidation number will be +1

note - does not work for halogens in complex ions, the charge in a complex ion represents a missing element

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7
Q

rule 5

A

oxygen has an oxidation number of -2. except in the compound OF₂ it is +2 as F is more electronegative element than O and in the peroxide ion [O₂²⁻] where O has an oxidation number of -1

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8
Q

rule 6

A

Hydrogen has an oxidation number of +1 in its compounds: except when it is combined to an element less electronegative than it because then it is a hydride and has an oxidation number of -1

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9
Q

hydride molecule

A

in a hydride molecule, there are two elements in the molecule and hydrogen is written last. Also the metal element will have a lower electronegativity value than hydrogen

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10
Q

rule 7

A

In complex ions, i.e. where more than one element is in the ion, the oxidation numbers must add up to the charge on the ion

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11
Q

rule 8

A

in neutral molecules, the oxidation numbers must add up to zero

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12
Q

oxidation

A

addition of oxygen

loss of electrons

increase in oxidation number

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13
Q

reduction

A

loss of oxygen

gain of electrons

decrease in oxidation number

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14
Q

oxidising agent

A

causes oxidation and is itself reduced

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15
Q

reducing agent

A

causes reduction and is itself oxidised

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16
Q

oxidation + reduction

A
  • Before discovery of electrons, reactions that involved addition of oxygen to a substance described as oxidation reactions
  • Removing oxygen from substance described as reduction reaction as mass of substance got smaller due to oxygen being removed
  • After discovery of electron, scientists examined reactions more closely + noticed many chem reactions involve transfer of electrons
  • One substance lost electrons while other gained
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17
Q

eg of oxidation reactions (before discovery of electrons, reactions that involved addition of oxygen …..) [continued point]

A
  • Burning of coal, which is carbon to produce carbon dioxide was an oxidation
  • The rusting of iron to produce iron oxide
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18
Q

oxidation/reduction electron transfer

A
  • Oxidation is when an element loses electrons

- reduction is when an element gains electrons

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19
Q

electron transfer

A
  • both oxidation + reduction occur at same time. If one element loses electrons, another gains electrons
  • called REDOX reactions
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20
Q

eg of electron transfer

A

in hback

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21
Q

The Electrochemical series

A

list of elements in order of their standard electrode potential

  • top: readily lose electrons
  • bottom: unreactive
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22
Q

reactivity

A

how easily they lose electrons

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23
Q

reactivity series of metals

A

in hback

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24
Q

electrolysis

A

the use of electricity to bring about a chemical reaction in an electrolyte

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25
electrolyte
a compound which when molten or dissolved in water will conduct an electric current
26
conduction of electricity
the conduction of electricity is due to the presence of ions
27
electrodes
-the two rods that clip into the electrolyte and make electrical contact with it are called the electrodes
28
two types of electrodes
- Inert | - active
29
inert electrode
- do not react w/ electrolye - graphite (carbon) - platinum
30
active electrode
- react with electrode - copper - iron
31
one elecrode is __ and the other is __
one is positive, other is negative
32
cathode
negative
33
anode
positive
34
to remember which is positive and negative
CNAP
35
electrodes proceedure
Battery 'pumps' electrons to the negative electrode where they are gained by some species + carried to positive electrode where they are lost
36
what happens at node?
oxidation
37
what happens at cathode?
reduction
38
to investigate what happens when an electrical current passes thru a soln of potassium iodide using inert electrodes
in hback
39
to investigate the electrolysis of acidified water using inert electrodes
- electric current passed thru acidified water - dilute sulfuric acid used - overall: H₂O + H₂ --> 1/2 O₂ - cathode - anode - SO4²⁻ not oxidised at anode (positive) bc they are v stable + loss of electrons would require a lot of energy - This is why HCl not used - Cl⁻ ions would lose electrons at anode + form chlorine gas Cl₂ - Conc of sulfuric acid remains constant (For every H⁺ ion reduced at cathode, one is produced at anode)
40
to investigate the electrolysis of acidified water using inert electrodes: cathode
- cathode (-): 2H⁺ + 2e⁻ --> H₂ (Reduction) - H⁺ ion from sulfuric acid are attracted to the negative electrode - H⁺ ion gains an electron + becomes H atom. 2H atoms bond to form H₂ gas
41
to investigate the electrolysis of acidified water using inert electrodes: anode
- Anode (-): H₂O --> H₂ + 1/2 O₂ + 2e⁻ - Electrons removed from the H₂O - Cause it to break into H⁺ and oygen gas
42
To investigate the electrolysis of sodium sulfate solution using inert electrodes
- Cathide - Anode - Na⁺ ions and SO₄²⁻ ions do note take part in reaction + are v stable
43
To investigate the electrolysis of sodium sulfate solution using inert electrodes: cathode
- Cathode (-): 2H₂O + 2e⁻ --> H₂ + 2OH⁻ (reduction) | - Turns blue in Universal Indicator as it produces OH⁻ ions (base)
44
To investigate the electrolysis of sodium sulfate solution using inert electrodes: anode
- Anode (+): 2H₂O --> 4H₂ + O₂ + 4e⁻ (oxidation) | - Turns red in Universal Indicator as it produces H⁺ ions (acid)
45
voltameter diagram
hback
46
To investigate the electrolysis of copper sulfate solution using copper electrodes
- Cathode | - Anode
47
To investigate the electrolysis of copper sulfate solution using copper electrodes: cathode
- Cathode (-): Cu²⁺ + 2e⁻ --> Cu (reduction) - Cu²⁺ ion in soln attracted to negative electrode - Cu²⁺ accepts 2e⁻ from copper electrode + forms copper metal which is plated onto copper electrode
48
To investigate the electrolysis of copper sulfate solution using copper electrodes: anode
- Anode (+): Cu --> Cu²⁺ + 2e⁻ (oxidation) - Cu atoms in positive electrode lose 2 electrons + become Cu²⁺ ions - Need tp replace Cu²⁺ ions lost in soln at negative electrode - Positive electrode is 'eaten away' while negative electrode 'gains weight'
49
Electroplating
-Process where electrolysis is used to put a layer of one metal onto another
50
Electroplating: Object placement
-Object to be plated at cathode (-) (reduction - ions gain electrons + form metals) -electrolyte must contain ions of the metal being plated -Metal being plated is placed at anode (+) (oxidised + goes into soln) -Plating silver
51
galvanic cell
a cell in which a chemical reaction results in an electric current
52
galvanic cell proceedure(?)
- Metals at each electrode need to have diff electrode potentials - One needs to be higher on electrochemical series than other so there is a pull on electrons - Wont work if metals at each electrode are same (no pull) - Electrons fly thru wire from one electrode to other (towards electrode with greater pull) - In this cell Zinc more reactive than copper + will lose electrons (oxidation) - Zinc becomes positively charged (anode) - Electrons fly to copper electrode making it negative (cathode) - Reduction occur at cathode + copper is plated onto electrode
53
galvanic cell proceedure(?) - zinc becoming positively charged
Zn --> Zn²⁺ + 2e⁻
54
galvanic cell proceedure(?) - reduction at end/plating at end
Cu²⁺ + 2e⁻ --> Cu
55
define reduction in terms of electron transfer
gain of electrons
56
define reduction in terms of change in oxidation number
reduction/decrease in oxidation number
57
define oxidation in terms of electron transfer
loss of electrons
58
define oxidation in terms of change in oxidation number
increase in oxidation number
59
colour change observed at positive electrode during electrolysis of potassium iodide
colour of the iodine produced
60
how does the oxidation number of the oxidising agent change during a redox reaction?
it decreases
61
why does the oxidising ability of the halogens decrease down the group?
- increasing atomic radius - increase in shielding - decrease in electronegativity (attraction for electrons)
62
state and explain the oxidation number of oxygen in the compound OF₂
oxygen is less electronegative fluorine is more electronegative
63
what is observed when chlorine gas is bubbled into an aqueous solution of sodium bromide?
turns red-brown
64
a solution of acidified water is electrolysed by passing an electric current through it using inert electrodes. Which electrode, A (left) or B (right) does oxidation occur Which species is oxidised?
A / positive electrode / anode H₂O (water)
65
keep in mind when balancing through oxidation
1: balance atoms first 2: balance oxidation numbers, keeping mole numbers in mind.. Eg, if H+ becomes 4H+ then the oxidation number is +4 3: check to see if atoms need to be balanced again after balancing oxidation numbers
66
electrolysis equations
learn off, in hardback
67
why decolorising time became shorter when second portion of potassium manganate solution was added to ethanedioic acid
catalyst produced by reaction + more catalyst produced as more KMnO4 reacted
68
tablet experiment calculations - finding mass of iron in 1 tablet
- When you get the molarity, convert it from mol/L to mol/volume of soln (eg. 250 cm3), - then use that value in n=m/mr to get the number of moles for 4 tablets - then divide by 4 to get mass for one tablet

LC Chemistry (26 decks)

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