5.3 Transition elements Flashcards
d-block element definition
element that has atoms with highest energy level electron in d-orbital
transition element definition
forms stable ion with incomplete d-subshell
why scandium isn’t transition element
loses 3 electrons as a 3+ ion
no d-sub shell at all
why zinc isn’t transition element
loses 2 electrons as a 2+ ion from s-sub shell
d-sub shell is complete
how d-block elements lose electrons
always lose 4s subshell electrons first
chromium electron configuration
4s contains one electron
3d orbitals all only contain 1 electron
copper electron configuration
all 3d orbitals full
4s only has one
why chromium and copper have weird electron config
thought to reduce repulsion of electrons
why transition metals can act as catalyst
can gain or lose electrons in d-subshell easily
easily transfer electrons to speed up reactions
complex ion definition
metal ion bonded to one or more ligands by coordinate bonds
coordinate bond definition
dative covalent bond
ligand definition
molecule or ion that can donate a pair of electrons to the transition metal ion to form a coordinate bond
coordination number definition
total number of coordinate bonds formed in the complex ion
monodentate ligand definition
only 1 atom in molecule/ion will donate the lone pair to the metal
bidentate ligands definition
2 atoms in the molecule/ion will donate the lone pair to the metal
monodentate ligands examples
water ammonia chloride cyanide hydroxide
bidentate ligand examples
1,2-diaminoethane ethanedioate ion (oxalate ion)
what complex ions can show cis-trans isomerism
square planar 4-coordinate complexes
6-coordinate complexes
what complex ions can show optical isomerism
tetrahedral 4-coordinate complexes
6-coordinate complexes
cis-trans isomerism in 6-coordinate complexes and square planar 4-coordinate complexes
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
what optical isomers do to polarised light
rotate plane-polarised light clockwise or anti-clockwise
chiral molecule definition
has a non-super imposable mirror image
optical isomerism requirements for tetrahedral
all groups need to be different
ligand substitution
reaction where one ligand in a complex ion is replaced by another ligand
precipitation reaction definition
2 aqueous solutions containing ions react together to form an insoluble ionic solid (precipitate)
obs when Cu2+(aq) + NaOH(aq)
blue solution to blue precipitate
insoluble in excess NaOH
Cu2+ (aq) + 2OH- (aq) -> Cu(OH)2 (s)
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
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)
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)
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)
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)
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)
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)
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)
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)
CuSO4 dissolved in water forms
Cu^2+ + 6H2O- -> [Cu(H2O)6]^2+
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)
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)
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
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
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)
how [Cr(H2O)5(SO4)]+ is formed
chromium (III) sulfate Cr2(SO4)3 dissolved in water
[Cr(H2O)5(SO4)]+ formed (green solution)
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)
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
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
reduction definition
gain of electrons
decrease in oxidation number
oxidation definition
loss of electrons
increase in oxidation number
oxidising agent definition
accepts pair of electrons from species being oxidised
reducing agent definition
donates pair of electrons to species being reduced
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)
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
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)
voltaic cell definition
type of electrochemical cell which converts chemical energy to electrical energy
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)
half cell definition
contains chemical species present in redox half equation
metal/metal ion half cell definition
consists of metal rod dipped in solution of aqueous ions
ion/ion half cell definition
contains ions of same elements in different oxidation states
why platinum used as an electrode in ion/ion half cell
inert
no metal to transfer electrons in or out the half cell
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
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
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)
E ⦵ (cell) formula
E ⦵(cell) = E ⦵(positive electrode) - E ⦵(negative electrode)
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
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
sigma bond definition
direct head-on overlap of orbitals between atoms
pi bond definition
sideways overlap of p-orbitals
primary cells features
not rechargeable (one time use)
alkaline-based
made up of Zn, MgO and KOH electrolyte
secondary cell features
rechargeable
lead-acid batteries (used in cars)
NiCd (used in radios, torches)
Li-ion and Li-ion polymer cells (modern appliances)
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
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
disadv of primary cells
low current
less environmentally friendly (harder to recycle)
large batteries less cost effective
one-time use (large amount of waste)
adv of secondary cells
rechargeable
chemicals regenerated
Li-ion is lightweight
more cost-efficient over time
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
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
disadv of fuel cells
hydrogen hard to store expensive less durable (not as long-lasting) no hydrogen fueling stations difficult to make batteries
anodes and cathodes in electrochemical cells
opposite
anode is negative
cathode is positive
Ecell of acid and alkali hydrogen fuel cells
both produce +1.23V
alkali fuel cell half equations and voltages
2H2O + 2e- ⇌ H2 + OH-
E = -0.83V (anode)
1/2O2 + H2O + 2e- ⇌ 2OH-
E = +0.40V (cathode)
acid fuel cell half equations
2H+ + 2e- ⇌ H2
E = 0.00 V (anode)
1/2O2 + 2H+ + 2e- ⇌ H2O
E = +1.23V (cathode)
CrO4^- colour
yellow
