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A-Level Chemistry > 5.3 Transition elements > Flashcards

5.3 Transition elements Flashcards

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1
Q

d-block element definition

A

element that has atoms with highest energy level electron in d-orbital

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2
Q

transition element definition

A

forms stable ion with incomplete d-subshell

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3
Q

why scandium isn’t transition element

A

loses 3 electrons as a 3+ ion

no d-sub shell at all

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4
Q

why zinc isn’t transition element

A

loses 2 electrons as a 2+ ion from s-sub shell

d-sub shell is complete

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5
Q

how d-block elements lose electrons

A

always lose 4s subshell electrons first

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6
Q

chromium electron configuration

A

4s contains one electron

3d orbitals all only contain 1 electron

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7
Q

copper electron configuration

A

all 3d orbitals full

4s only has one

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8
Q

why chromium and copper have weird electron config

A

thought to reduce repulsion of electrons

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9
Q

why transition metals can act as catalyst

A

can gain or lose electrons in d-subshell easily

easily transfer electrons to speed up reactions

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10
Q

complex ion definition

A

metal ion bonded to one or more ligands by coordinate bonds

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11
Q

coordinate bond definition

A

dative covalent bond

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12
Q

ligand definition

A

molecule or ion that can donate a pair of electrons to the transition metal ion to form a coordinate bond

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13
Q

coordination number definition

A

total number of coordinate bonds formed in the complex ion

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14
Q

monodentate ligand definition

A

only 1 atom in molecule/ion will donate the lone pair to the metal

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15
Q

bidentate ligands definition

A

2 atoms in the molecule/ion will donate the lone pair to the metal

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16
Q

monodentate ligands examples

A 
water
ammonia
chloride
cyanide
hydroxide
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17
Q

bidentate ligand examples

A 
1,2-diaminoethane 
ethanedioate ion (oxalate ion)
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18
Q

what complex ions can show cis-trans isomerism

A

square planar 4-coordinate complexes

6-coordinate complexes

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19
Q

what complex ions can show optical isomerism

A

tetrahedral 4-coordinate complexes

6-coordinate complexes

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20
Q

cis-trans isomerism in 6-coordinate complexes and square planar 4-coordinate complexes

A

ligand of interest on same side (adjacent to each other) so 90° bond angle = cis
ligand of interest on different sides so 180° bond angle = trans

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21
Q

what optical isomers do to polarised light

A

rotate plane-polarised light clockwise or anti-clockwise

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22
Q

chiral molecule definition

A

has a non-super imposable mirror image

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23
Q

optical isomerism requirements for tetrahedral

A

all groups need to be different

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24
Q

ligand substitution

A

reaction where one ligand in a complex ion is replaced by another ligand

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25
Q

precipitation reaction definition

A

2 aqueous solutions containing ions react together to form an insoluble ionic solid (precipitate)

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26
Q

obs when Cu2+(aq) + NaOH(aq)

A

blue solution to blue precipitate
insoluble in excess NaOH
Cu2+ (aq) + 2OH- (aq) -> Cu(OH)2 (s)

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27
Q

obs when Cu2+(aq) + NH3(aq)

A

blue solution to blue precipitate to dark blue solution
soluble in excess NH3
Cu2+ (aq) + 2OH- (aq) -> Cu(OH)2 (s)
in excess NH3(aq):
[Cu(H2O)6]2+ (aq) + 4NH3(aq) -> [Cu(NH3)4(H2O)2]2+ (aq) + 4H2O (l)
Cu(OH2) dissolves in excess ammonia

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28
Q

obs when Fe2+(aq) + NaOH(aq)

A

pale green solution to green precipitate
insoluble in excess NaOH
precipitate turns brown if exposed to air
Fe2+ (aq) + 2OH- (aq) -> Fe(OH)2 (s)
in air: Fe(OH)2 (s) -> Fe(OH)3 (s)

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29
Q

obs when Fe2+(aq) + NH3(aq)

A

pale green solution to green precipitate
insoluble in excess NH3
Fe2+ (aq) + 2OH- (aq) -> Fe(OH)2 (s)
in air: Fe(OH)2 (s) -> Fe(OH)3 (s)

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30
Q

obs when Fe3+(aq) + NaOH(aq)

A

pale yellow solution to orange/brown precipitate
insoluble in excess NaOH
Fe3+ (aq) + 3OH- (aq) -> Fe(OH)3 (s)

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31
Q

obs when Fe3+(aq) + NH3(aq)

A

pale yellow solution to orange/brown precipitate
insoluble in excess NH3
Fe3+ (aq) + 3OH- (aq) -> Fe(OH)3 (s)

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32
Q

obs when Mn2+(aq) + NaOH(aq)

A

pale pink solution to light brown precipitate that darkens on standing in air
insoluble in excess NaOH
Mn2+ (aq) + 2OH-(aq) -> Mn(OH)2 (s)

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33
Q

obs when Mn2+ (aq) + NH3(aq)

A

pale pink solution to light brown precipitate, darkens standing on air
insoluble in excess NH3(aq)
Mn2+ (aq) + 2OH- (aq) -> Mn(OH)2 (aq)

34
Q

obs when Cr3+ (aq) + NaOH(aq)

A

violet solution reacts to grey-green precipitate
soluble in excess NaOH(aq) to form dark green solution
Cr3+(aq) + 3OH- (aq) -> Cr(OH)3 (s)
in excess:
Cr(OH)3 (s) + 3OH- (aq) -> [Cr(OH)6]3- (aq)

35
Q

obs when Cr3+ + NH3(aq)

A

violet solution to grey-green precipitate
soluble in excess ammonia to form purple solution
Cr3+ (aq) + 3OH- (aq) -> Cr(OH)3 (s)
in excess NH3:
Cr3+ (aq)+ 6NH3(aq) -> [Cr(NH3)6]3+ (aq)

36
Q

CuSO4 dissolved in water forms

A

Cu^2+ + 6H2O- -> [Cu(H2O)6]^2+

37
Q

ligand substitution of [Cu(H2O)6]2+ with ammonia

A

pale blue solution to dark blue solution

[Cu(H2O)6]2+ (aq) + 4NH3 (aq) -> [Cu(NH3)4(H2O)2]2+ (aq) + 4H2O (l)

38
Q

ligand substitution of [Cu(H2O)6]2+ with chloride ions

A

pale blue solution to yellow solution
excess conc. HCl(aq)
equilibrium reaction
[Cu(H2O)6]2+ (aq) + 4Cl-(aq) ⇌ [CuCl4]2- (aq) + 6H2O(l)

39
Q

why green intermediary made in ligand substitution of [Cu(H2O)6]2+ and Cl-

A

intermediate green solution formed is yellow solution mixing with blue solution as reaction proceeds

40
Q

why change in coordination number in ligand substitution of [Cu(H2O)6]2+ and Cl-

A

octahedral (6) to tetrahedral (4)

chloride ligands larger than water ligands so less can fit around central Cu2+ ion

41
Q

how [Cr(H2O)6]3+ is formed

A

KCr(SO4)2• 12H2O (chromium (III) potassium sulfate) dissolved in water
forms [Cr(H2O)6]3+ ions (violet solution)

42
Q

how [Cr(H2O)5(SO4)]+ is formed

A

chromium (III) sulfate Cr2(SO4)3 dissolved in water

[Cr(H2O)5(SO4)]+ formed (green solution)

43
Q

reaction of [Cr(H2O)6]3+ with ammonia

A

initially grey-green precipitate formed, Cr(OH)3 (s)
then dissolves in excess ammonia
forms [Cr(NH3)6]3+ (aq) (purple solution)

44
Q

importance of ligand substitution of Fe2+ in haemoglobin

A

Fe2+ allows haemoglobin to bind to O2 gas at high oxygen pressure to form oxyhaemoglobin
releases oxygen at low oxygen pressure (respiring tissue)
allows haem. to bind to CO2 at respiring tissue and releases it at the lungs to be exhaled

45
Q

carbon monoxide and haemoglobin

A

when inhaled, ligand substitution with CO and O2 in haemoglobin to form carboxyhaemoglobin
prevents large proportion of haemoglobin to carry oxygen
CO bond stronger than O2 bond
can lead to death

46
Q

reduction definition

A

gain of electrons

decrease in oxidation number

47
Q

oxidation definition

A

loss of electrons

increase in oxidation number

48
Q

oxidising agent definition

A

accepts pair of electrons from species being oxidised

49
Q

reducing agent definition

A

donates pair of electrons to species being reduced

50
Q

manganate titration method

A

standard solution of potassium manganate (VII) added to burette
add measured volume of solution being analysed to conical flask using pipette
add excess of dilute H2SO4(aq) to conical flask
during titration, manganate solution decolourised as it reacts
end point is when first permanent pink colour occurs (no indicator required)

51
Q

how to read meniscus during manganate titration

A

KMnO4(aq) is deep purple so hard to see bottom of meniscus

burette readings taken from top of meniscus rather than bottom

52
Q

iodine/thiosulfate titration method

A

add standard solution of Na2S2O3 to burette
prepare solution of oxidising agent to be analysed
add this solution to conical flask using pipette
add excess potassium iodide to conical flask
oxidising agent reacts with I- to produce iodine (turn solution yellow/brown)
titrate this solution with NaS2O3
iodine reduced back into I- ions
starch added to see clear endpoint (deep black-blue when iodine present, clear straw colour when it isnt)

53
Q

voltaic cell definition

A

type of electrochemical cell which converts chemical energy to electrical energy

54
Q

how electrode potentials work (Mg example)

A

electrode potentials (EPs) compare ease of metal to give up electrons to form positive hydrated ions

Mg loses 2e- (becomes Mg2+)
Mg2+ attracted to negative strip
picks up electrons again to become Mg(s)
Mg2+(aq) + e- <=> Mg(s) 
more reactive so equilibrium lies to LHS (as more Mg2+(aq) than Mg(s) is formed)
55
Q

half cell definition

A

contains chemical species present in redox half equation

56
Q

metal/metal ion half cell definition

A

consists of metal rod dipped in solution of aqueous ions

57
Q

ion/ion half cell definition

A

contains ions of same elements in different oxidation states

58
Q

why platinum used as an electrode in ion/ion half cell

A

inert

no metal to transfer electrons in or out the half cell

59
Q

standard electrode potential definition

A 
E ⦵
electromotive force (emf) of a half cell compared with a standard hydrogen half cell measured at 298K with solution concentrarions of 1.0 mol dm^-3 and a gas pressure of 100kPa
60
Q

salt bridge function when connecting half cells to make a cell

A

allows ions to flow but contains a solution that doesnt react with the half cell solution (e.g. filter paper soaked in KNO3

61
Q

positive or negative in an operating cell

A

electrode with more reactive metal(element) loses more electrons than it gains so is more negative (therefore oxidised and stronger reducing agent)
electrode with less reactive metal(element) gains more electrons than it loses so is more positive (therefore reduced and stronger oxidising agent)

62
Q

E ⦵ (cell) formula

A

E ⦵(cell) = E ⦵(positive electrode) - E ⦵(negative electrode)

63
Q

feasibility of reactions using electrode potentials

A

the more E°(cell) is greater than 0, the more feasible the reaction
also look if the species required for reaction to continue are available

64
Q

limitations of using E° to determine feasibility of reactions

A 
reaction rate (reaction may not occur due to high activation energy, E° gives no indication of this)  
concentration (E° measured in 1 mol dm^-3, many reactions take place in more or less concentrated solutions, E value and overall E(cell) will change
actual conditions carried out may different so E° no longer real values
many reactions take place that aren’t aqueous
65
Q

sigma bond definition

A

direct head-on overlap of orbitals between atoms

66
Q

pi bond definition

A

sideways overlap of p-orbitals

67
Q

primary cells features

A

not rechargeable (one time use)
alkaline-based
made up of Zn, MgO and KOH electrolyte

68
Q

secondary cell features

A

rechargeable
lead-acid batteries (used in cars)
NiCd (used in radios, torches)
Li-ion and Li-ion polymer cells (modern appliances)

69
Q

fuel cell features

A

uses energy from reaction of a fuel with oxygen to create a voltage
hydrogen most common as produces no CO2 by-product

70
Q

adv. of primary cells

A 
longer shelf-life and last longer for the charge they have
easier to replace in the field
cheaper
don’t need to be charged
come in various sizes
71
Q

disadv of primary cells

A

low current
less environmentally friendly (harder to recycle)
large batteries less cost effective
one-time use (large amount of waste)

72
Q

adv of secondary cells

A

rechargeable
chemicals regenerated
Li-ion is lightweight
more cost-efficient over time

73
Q

disadv of secondary cells

A 
expensive (initially)
poorer charge retention (over time voltage reduces)
can take long time to recharge
unstable at high temperatures 
difficult to recycle
74
Q

adv of fuel cells

A 
low pollutant
high efficiency
don’t have to be recharged as long as H2 and O2 supplied can operate continuously
no CO2 produced for hydrogen fuel cells
removed reliance on fossil fuels
75
Q

disadv of fuel cells

A 
hydrogen hard to store
expensive
less durable (not as long-lasting)
no hydrogen fueling stations
difficult to make batteries
76
Q

anodes and cathodes in electrochemical cells

A

opposite
anode is negative
cathode is positive

77
Q

Ecell of acid and alkali hydrogen fuel cells

A

both produce +1.23V

78
Q

alkali fuel cell half equations and voltages

A

2H2O + 2e- ⇌ H2 + OH-
E = -0.83V (anode)

1/2O2 + H2O + 2e- ⇌ 2OH-
E = +0.40V (cathode)

79
Q

acid fuel cell half equations

A

2H+ + 2e- ⇌ H2
E = 0.00 V (anode)

1/2O2 + 2H+ + 2e- ⇌ H2O
E = +1.23V (cathode)

80
Q

CrO4^- colour

A

yellow

A-Level Chemistry (12 decks)

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