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LC Chemistry > Oxidation Numbers > Flashcards

Oxidation Numbers Flashcards

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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
Q

electrolyte

A

a compound which when molten or dissolved in water will conduct an electric current

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26
Q

conduction of electricity

A

the conduction of electricity is due to the presence of ions

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27
Q

electrodes

A

-the two rods that clip into the electrolyte and make electrical contact with it are called the electrodes

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28
Q

two types of electrodes

A
  • Inert

- active

29
Q

inert electrode

A
  • do not react w/ electrolye
  • graphite (carbon)
  • platinum
30
Q

active electrode

A
  • react with electrode
  • copper
  • iron
31
Q

one elecrode is __ and the other is __

A

one is positive, other is negative

32
Q

cathode

A

negative

33
Q

anode

A

positive

34
Q

to remember which is positive and negative

A

CNAP

35
Q

electrodes proceedure

A

Battery ‘pumps’ electrons to the negative electrode where they are gained by some species + carried to positive electrode where they are lost

36
Q

what happens at node?

A

oxidation

37
Q

what happens at cathode?

A

reduction

38
Q

to investigate what happens when an electrical current passes thru a soln of potassium iodide using inert electrodes

A

in hback

39
Q

to investigate the electrolysis of acidified water using inert electrodes

A
  • 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
Q

to investigate the electrolysis of acidified water using inert electrodes: cathode

A
  • 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
Q

to investigate the electrolysis of acidified water using inert electrodes: anode

A
  • Anode (-): H₂O –> H₂ + 1/2 O₂ + 2e⁻
  • Electrons removed from the H₂O
  • Cause it to break into H⁺ and oygen gas
42
Q

To investigate the electrolysis of sodium sulfate solution using inert electrodes

A
  • Cathide
  • Anode
  • Na⁺ ions and SO₄²⁻ ions do note take part in reaction + are v stable
43
Q

To investigate the electrolysis of sodium sulfate solution using inert electrodes: cathode

A
  • Cathode (-): 2H₂O + 2e⁻ –> H₂ + 2OH⁻ (reduction)

- Turns blue in Universal Indicator as it produces OH⁻ ions (base)

44
Q

To investigate the electrolysis of sodium sulfate solution using inert electrodes: anode

A
  • Anode (+): 2H₂O –> 4H₂ + O₂ + 4e⁻ (oxidation)

- Turns red in Universal Indicator as it produces H⁺ ions (acid)

45
Q

voltameter diagram

A

hback

46
Q

To investigate the electrolysis of copper sulfate solution using copper electrodes

A
  • Cathode

- Anode

47
Q

To investigate the electrolysis of copper sulfate solution using copper electrodes: cathode

A
  • 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
Q

To investigate the electrolysis of copper sulfate solution using copper electrodes: anode

A
  • 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
Q

Electroplating

A

-Process where electrolysis is used to put a layer of one metal onto another

50
Q

Electroplating: Object placement

A

-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
Q

galvanic cell

A

a cell in which a chemical reaction results in an electric current

52
Q

galvanic cell proceedure(?)

A
  • 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
Q

galvanic cell proceedure(?) - zinc becoming positively charged

A

Zn –> Zn²⁺ + 2e⁻

54
Q

galvanic cell proceedure(?) - reduction at end/plating at end

A

Cu²⁺ + 2e⁻ –> Cu

55
Q

define reduction in terms of electron transfer

A

gain of electrons

56
Q

define reduction in terms of change in oxidation number

A

reduction/decrease in oxidation number

57
Q

define oxidation in terms of electron transfer

A

loss of electrons

58
Q

define oxidation in terms of change in oxidation number

A

increase in oxidation number

59
Q

colour change observed at positive electrode during electrolysis of potassium iodide

A

colour of the iodine produced

60
Q

how does the oxidation number of the oxidising agent change during a redox reaction?

A

it decreases

61
Q

why does the oxidising ability of the halogens decrease down the group?

A
  • increasing atomic radius
  • increase in shielding
  • decrease in electronegativity (attraction for electrons)
62
Q

state and explain the oxidation number of oxygen in the compound OF₂

A

oxygen is less electronegative

fluorine is more electronegative

63
Q

what is observed when chlorine gas is bubbled into an aqueous solution of sodium bromide?

A

turns red-brown

64
Q

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

A / positive electrode / anode

H₂O (water)

65
Q

keep in mind when balancing through oxidation

A

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
Q

electrolysis equations

A

learn off, in hardback

67
Q

why decolorising time became shorter when second portion of potassium manganate solution was added to ethanedioic acid

A

catalyst produced by reaction + more catalyst produced as more KMnO4 reacted

68
Q

tablet experiment calculations - finding mass of iron in 1 tablet

A
  • 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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