5.3 Transition elements Flashcards
1
Q
d-block element definition
A
element that has atoms with highest energy level electron in d-orbital
2
Q
transition element definition
A
forms stable ion with incomplete d-subshell
3
Q
why scandium isn’t transition element
A
loses 3 electrons as a 3+ ion
no d-sub shell at all
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
5
Q
how d-block elements lose electrons
A
always lose 4s subshell electrons first
6
Q
chromium electron configuration
A
4s contains one electron
3d orbitals all only contain 1 electron
7
Q
copper electron configuration
A
all 3d orbitals full
4s only has one
8
Q
why chromium and copper have weird electron config
A
thought to reduce repulsion of electrons
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
10
Q
complex ion definition
A
metal ion bonded to one or more ligands by coordinate bonds
11
Q
coordinate bond definition
A
dative covalent bond
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
13
Q
coordination number definition
A
total number of coordinate bonds formed in the complex ion
14
Q
monodentate ligand definition
A
only 1 atom in molecule/ion will donate the lone pair to the metal
15
Q
bidentate ligands definition
A
2 atoms in the molecule/ion will donate the lone pair to the metal
16
Q
monodentate ligands examples
A
water ammonia chloride cyanide hydroxide
17
Q
bidentate ligand examples
A
1,2-diaminoethane ethanedioate ion (oxalate ion)
18
Q
what complex ions can show cis-trans isomerism
A
square planar 4-coordinate complexes
6-coordinate complexes
19
Q
what complex ions can show optical isomerism
A
tetrahedral 4-coordinate complexes
6-coordinate complexes
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
21
Q
what optical isomers do to polarised light
A
rotate plane-polarised light clockwise or anti-clockwise
22
Q
chiral molecule definition
A
has a non-super imposable mirror image
23
Q
optical isomerism requirements for tetrahedral
A
all groups need to be different
24
Q
ligand substitution
A
reaction where one ligand in a complex ion is replaced by another ligand
25
precipitation reaction definition
2 aqueous solutions containing ions react together to form an insoluble ionic solid (precipitate)
26
obs when Cu2+(aq) + NaOH(aq)
blue solution to blue precipitate
insoluble in excess NaOH
Cu2+ (aq) + 2OH- (aq) -> Cu(OH)2 (s)
27
obs when Cu2+(aq) + NH3(aq)
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
28
obs when Fe2+(aq) + NaOH(aq)
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)
29
obs when Fe2+(aq) + NH3(aq)
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)
30
obs when Fe3+(aq) + NaOH(aq)
pale yellow solution to orange/brown precipitate
insoluble in excess NaOH
Fe3+ (aq) + 3OH- (aq) -> Fe(OH)3 (s)
31
obs when Fe3+(aq) + NH3(aq)
pale yellow solution to orange/brown precipitate
insoluble in excess NH3
Fe3+ (aq) + 3OH- (aq) -> Fe(OH)3 (s)
32
obs when Mn2+(aq) + NaOH(aq)
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)
33
obs when Mn2+ (aq) + NH3(aq)
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
obs when Cr3+ (aq) + NaOH(aq)
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
obs when Cr3+ + NH3(aq)
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
CuSO4 dissolved in water forms
Cu^2+ + 6H2O- -> [Cu(H2O)6]^2+
37
ligand substitution of [Cu(H2O)6]2+ with ammonia
pale blue solution to dark blue solution
| [Cu(H2O)6]2+ (aq) + 4NH3 (aq) -> [Cu(NH3)4(H2O)2]2+ (aq) + 4H2O (l)
38
ligand substitution of [Cu(H2O)6]2+ with chloride ions
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
why green intermediary made in ligand substitution of [Cu(H2O)6]2+ and Cl-
intermediate green solution formed is yellow solution mixing with blue solution as reaction proceeds
40
why change in coordination number in ligand substitution of [Cu(H2O)6]2+ and Cl-
octahedral (6) to tetrahedral (4)
| chloride ligands larger than water ligands so less can fit around central Cu2+ ion
41
how [Cr(H2O)6]3+ is formed
KCr(SO4)2• 12H2O (chromium (III) potassium sulfate) dissolved in water
forms [Cr(H2O)6]3+ ions (violet solution)
42
how [Cr(H2O)5(SO4)]+ is formed
chromium (III) sulfate Cr2(SO4)3 dissolved in water
| [Cr(H2O)5(SO4)]+ formed (green solution)
43
reaction of [Cr(H2O)6]3+ with ammonia
initially grey-green precipitate formed, Cr(OH)3 (s)
then dissolves in excess ammonia
forms [Cr(NH3)6]3+ (aq) (purple solution)
44
importance of ligand substitution of Fe2+ in haemoglobin
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
carbon monoxide and haemoglobin
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
reduction definition
gain of electrons
| decrease in oxidation number
47
oxidation definition
loss of electrons
| increase in oxidation number
48
oxidising agent definition
accepts pair of electrons from species being oxidised
49
reducing agent definition
donates pair of electrons to species being reduced
50
manganate titration method
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
how to read meniscus during manganate titration
KMnO4(aq) is deep purple so hard to see bottom of meniscus
| burette readings taken from top of meniscus rather than bottom
52
iodine/thiosulfate titration method
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
voltaic cell definition
type of electrochemical cell which converts chemical energy to electrical energy
54
how electrode potentials work (Mg example)
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
half cell definition
contains chemical species present in redox half equation
56
metal/metal ion half cell definition
consists of metal rod dipped in solution of aqueous ions
57
ion/ion half cell definition
contains ions of same elements in different oxidation states
58
why platinum used as an electrode in ion/ion half cell
inert
| no metal to transfer electrons in or out the half cell
59
standard electrode potential definition
```
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
salt bridge function when connecting half cells to make a cell
allows ions to flow but contains a solution that doesnt react with the half cell solution (e.g. filter paper soaked in KNO3
61
positive or negative in an operating cell
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
E ⦵ (cell) formula
E ⦵(cell) = E ⦵(positive electrode) - E ⦵(negative electrode)
63
feasibility of reactions using electrode potentials
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
limitations of using E° to determine feasibility of reactions
```
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
sigma bond definition
direct head-on overlap of orbitals between atoms
66
pi bond definition
sideways overlap of p-orbitals
67
primary cells features
not rechargeable (one time use)
alkaline-based
made up of Zn, MgO and KOH electrolyte
68
secondary cell features
rechargeable
lead-acid batteries (used in cars)
NiCd (used in radios, torches)
Li-ion and Li-ion polymer cells (modern appliances)
69
fuel cell features
uses energy from reaction of a fuel with oxygen to create a voltage
hydrogen most common as produces no CO2 by-product
70
adv. of primary cells
```
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
disadv of primary cells
low current
less environmentally friendly (harder to recycle)
large batteries less cost effective
one-time use (large amount of waste)
72
adv of secondary cells
rechargeable
chemicals regenerated
Li-ion is lightweight
more cost-efficient over time
73
disadv of secondary cells
```
expensive (initially)
poorer charge retention (over time voltage reduces)
can take long time to recharge
unstable at high temperatures
difficult to recycle
```
74
adv of fuel cells
```
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
disadv of fuel cells
```
hydrogen hard to store
expensive
less durable (not as long-lasting)
no hydrogen fueling stations
difficult to make batteries
```
76
anodes and cathodes in electrochemical cells
opposite
anode is negative
cathode is positive
77
Ecell of acid and alkali hydrogen fuel cells
both produce +1.23V
78
alkali fuel cell half equations and voltages
2H2O + 2e- ⇌ H2 + OH-
E = -0.83V (anode)
1/2O2 + H2O + 2e- ⇌ 2OH-
E = +0.40V (cathode)
79
acid fuel cell half equations
2H+ + 2e- ⇌ H2
E = 0.00 V (anode)
1/2O2 + 2H+ + 2e- ⇌ H2O
E = +1.23V (cathode)
80
CrO4^- colour
yellow
