Unit 5 Chemistry Notes

A2 Unit 5 Energetics, Redox and Inorganic Chemistry

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AQA A2-LEVEL

Student Guide to A2 Unit 5

Energetics, Redox and Inorganic Chemistry

See me in glorious at:

Shared Areas Chemistry Read Mr Lunds Classes A2 Chemistry Unit 5 All programs


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Enthalpy Changes and Hesss Law

enthalpy changeH = heat exchange at constant pressure (open container)

enthalpy change cannot be measured directly but can be determined experimentally

H varies with temperature and pressure so Standard Conditions (Ho298) are required:

temperature 298 K (25C i.e. nominal room temperature)

pressure 100kPa dont put 1 atmosphere in the exam!!! physical state at room temperature and most stable allotrope (e.g. graphite)

the enthalpy of formation of an element is by definition zero

here are two enthalpy changes that you learned in AS Chemistry:

the enthalpy of is the enthalpy change that occurs when one mole of

at 298K and 100 kPa (i.e. standard conditions)

formation (Hof,298) a compound is formed from its elements

combustion (oc,298) an element or compound is completely combusted in xs oxygen

Hesss Law: the enthalpy change for a reaction is dependent only on the initial and final states of the system and is independent of the route taken.

Hor,298 = Hof,298 (Products) - Hof,298 (Reactants)

Hor,298 = Hoc,298 (Reactants) - Hoc,298 (Products)

Enthalpy Definitions

Reactants

s

Products

Combustion

Products

2

1 Hesss Law states that 1 = 2

Reactants Products

Elements

2

1 Hesss Law states that 1 = 2

A MARK IN THE EXAM !!


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note that once again in all these definitions you specify both the quantity and type of the particles initially and finally



atomisation Hoat gaseous atoms are formed from an element in its standard state

note the differences between the enthalpies of formation, atomisation and bond dissociation as these are often used incorrectly in calculations (think about the values for Iodine)

1st ionisation energy Hoi(1) gaseous singly charged cations are formed from gaseous atoms

revise the trends (periods and groups) of first ionisation energies in addition to the changes in magnitude (shells) and equations for successive ionisation energies of the same element

1st electron affinity Hoea(1) gaseous singly charged anions are formed from gaseous atoms

you should be able to write an equation for the gain of an electron by any given atom/ion

the first electron affinity is exothermic as the vacancy in the subshell is subject to an attractive force from the nucleus that is not completely screened.

the second electron affinity is endothermic since an electron is being moved towards a negative ion and mutual repulsion must be overcome (subsequent additions are increasingly

endothermic as the size of the negative charge increases)

lattice enthalpy HoL an ionic lattice is formed from its gaseous ions

lattice dissociation enthalpy an ionic compound separates into gaseous ions

lattice enthalpy is higher if the charges are higher and/or the ions smaller as the electrostatic force of attraction is greater see 169 (note its the same as lattice energy for A-level use)

Lattice enthalpy and enthalpy of formation are not the same thing:

Lattice enthalpy starts from gaseous ions.

Enthalpy of formation starts from elements in their normal states.

lattice energy cannot be determined experimentally as it is impossible to react an exact amount of gaseous anions and cations and measure the energy released.

Note that these

have the same

numerical

value but

opposite signs

EXOTHERMIC

ENDOTHERMIC


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the value of the lattice energy can be derived theoretically using an electrostatic relationship between charge sizes and distances but this makes two incorrect assumptions:

1 the ions are point sources 2 the ions do not mutually interact such as to distort their shape i.e that they are purely ionic

and exhibit no degree of covalency

experimental values are higher than theoretically calculated values for HoL as there is in reality a degree of covalency in all ionic compounds as the electrostatic attraction of the

cations will distort the valence electron cloud of the anions

in most cases (small ions that are singly charged) there is a reasonable agreement between experimental and theoretical data (within 5%) i.e. the ionic model is adequate.

however, the deviation is more significant i.e. the degree of covalency increased if:

1 the cations are smaller 2 the anions are larger 3 the charge on the ions is larger

a small cation with a charge of 2+ or higher will present a highly polarising field to an anion the effect of which will be more

significant if that anion is relatively large

the anionic electron clouds will be distorted (polarised) towards the polarising field of the cation (this is not dissimilar to the tenuous

outer atmosphere of a red giant star becoming distorted in

association with a strong gravitational field)

this can lead to an overlap of electron clouds i.e. the compound is in effect covalent

Summary Questions Page 166 1 3 Page 174 2

A2 Chemistry (Nelson Thornes) AQA 163 166, 169, 174

Chemguide Lattice enthalpy, ionic structures

s-cool: Chemical Energetics


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Born-Haber Cycles

a Born-Haber Cycle for a simple ionic solid can be constructed using Hesss Law (including ionisation energy, electron affinity, lattice energy, atomisation energy)

Note that, unlike the energy cycle diagrams met at AS-level, these are energy level diagrams in

that they are drawn with up = endothermic and down = exothermic change

elements are referenced to zero

NaCl and MgCl2 are shown in your textbook

for oxides/sulphides the second electron affinity will be endothermic (i.e. back up)

but why not MgCl without a second (endothermic) ionisation energy or MgCl3 where a 3+ ion is forming the lattice giving a much higher (exothermic) lattice energy (see page 170)

ionisation energy and lattice energy are the two major contributory values

Enthalpies of Solution (Hosol)



solution Hosol solute is completely dissolved in water

hydration Hohyd isolated gaseous ions are hydrated

hydration involves ion-dipole electrostatic forces hence ionic solids can dissolve in polar solvents such as water

the payback of hydration enthalpy is not available in non-polar solvents rendering the process too endothermic overall to be possible

solubility is essentially a trade off between overcoming the enthalpy of lattice dissociation enthalpy (endothermic in this direction) and enthalpies of hydration (always exothermic)

solubility is feasible if this is an exothermic outcome overall (e.g. NaOH(aq))

although ENTROPY (see later) will also have a say if the enthalpy of solution is only slightly endothermic (e.g. NH4NO3(aq))

Hesss Law can be used to construct a solvation energy cycle (see 173)

Summary Questions Page 172 1

Page 174 1

How Science Works Pages 171 2 and Questions 1 - 4 Exam Style Questions Page 184 1

Page 248 5, 6 (read entropy for d(i))

A2 Chemistry (Nelson Thornes) AQA 164 173

Chemguide Born Haber lattice

s-cool: Chemical Energetics


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Free Energy and Entropy

it seems reasonable that exothermic reactions can occur spontaneously, after all the products are relatively more stable than the reactants (stronger bonds)

obviously the rate of the process will depend upon the kinetic barrier (EA) i.e. kinetic feasibility and they could be slow e.g. iron rusting or extremely slow e.g diamond into

graphite

spontaneously thus means it will happen under a given set of conditions but not necessarily in an instant and it could be almost infinitely slow

a small proportion of spontaneous reactions however are endothermic e.g. solvating NH4NO3(s)

at first sight this seems counter intuitive as endothermic reactions require a transfer of heat from their surrounding resulting in products that are relatively less stable (higher stored

chemical energy) in terms of enthalpy

so why doesnt NH4NO3(aq) spontaneously undissolve as this would be exothermic and also be kinetically more feasible relative to the forward reaction (lower EA)

why does it only go one way (assuming the solubility limit has not been exceeded)

clearly there must be another factor in the viability of a chemical (and physical) process that determines whether it will occur spontaneously or not

Entropy (S)

chemical and physical change are governed by the laws of probability

the degree of disorder (ways of distributing energy) within a system is called Entropy

the greater the degree of disorder in a system the greater the entropy e.g.

state changes

solvation

diffusion

complexity of molecules increases way to distribute energy hence greater entropy temperature entropy is increased if it is raised and visa-versa

the second law of thermodynamics states that all viable chemical and physical changes result in an increase in the TOTAL entropy in the Universe

hence all spontaneous reactions result in an increase in TOTAL entropy irrespective of whether they are exothermic or endothermic

an endothermic reaction can thus be spontaneous IF it results in an increased degree of randomness (of distribution of energy) in the universe


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S

heat energy flowing into a system will increase its entropy more available energy levels more available ways in which energy can be distributed between the particles

for an endothermic reaction lowering the temperature reduces entropy but the change from reactants to products could increase entropy (e.g. solvation).

IF there is a NET increase in entropy the reaction may be spontaneous

like enthalpy, So is specified for 298K and 100 kPa

unlike H it is possible to quantify S (units are JK-1mol-1) since it can be extrapolated from a value of zero at 0K

it is possible to determine the entropy change in the system (the chemicals reacting)

S(system) = S(products) - S(reactants)

note that the overall entropy may only change slightly if all the chemicals are in the same phase

if entropy must increase how can water condense spontaneously given that it achieves a more ordered state i.e. a decrease in entropy?

We must also consider the entropy change in the surroundings to complete the picture remember it is the TOTAL entropy change that concerns us

the total entropy change is expressed by:

Sototal = Sosystem + Sosurroundings

it is very difficult to calculate the entropy change in the surroundings

however, it can be determined from the enthalpy transferred from/to the system at a given temperature (since temperature effects the value of entropy)

T S o surroundings = -Hosystem

energy is transferred from the steam (system) to the surface that it condenses on (surroundings) increasing its entropy such that the overall entropy increases

the ligand displacement reactions that you will meet in transition metal chemistry can often be attributable to favourable entropy changes where the total number of particles increases

Look

!


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Gibbs free energy (G)

entropy and enthalpy can be combined in a quantity called Gibbs free energy (G)

G o is the standard free energy change (kjmol-1) and quantifies the enthalpy and entropy changes at a given temperature

for a reaction to be feasible it must be zero or negative i.e such that Sototal is positive

Go = - TSototal

some mathematical magic with the previous equations leads us to:

Go = Hosystem - TSosystem kjmol-1

G o effectively determines the feasibility of a reaction

reaction feasibility is thus determined by the relative magnitude of the enthalpy and entropy changess at a given temperature

temperature has a significant effect on reaction feasibility

a reaction first becomes feasible when the temperature has been raised to a value where G o = 0

this is particularly relevant to metal extraction where heating may be necessary to make the reaction feasible not just to increase the rate

a feasible reaction is not necessarily spontaneous as there may be a high activation energy involved

when G o = 0 an equilibrium will exist in a closed system

S Outcome

- + G always negative Feasible at ALL temperatures

+ - G always positive NOT feasible at ANY

temperature

- - IF > T S then reaction is feasible i.e. at LOWER

temperatures

+

+ IF T S > then reaction is feasible i.e. at HIGHER

temperatures

Remember to

divide entropy

by a 1000

since its in J

not kJ !!!


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Change of State

a system is at equilibrium (e.g. ice/water at 0oC) when:

G = 0

= T S

if heat goes into the system then the entropy term must increase for G to remain zero

i.e. it melts to a more disordered state

visa-versa is true

the entropy change for evaporation is significantly higher than for melting as the increase in randomness is more significant

Summary Questions Page 183 1 - 3

Exam Style Questions Page 185 2 4 Page 246 1, 2

A2 Chemistry (Nelson Thornes) AQA 178 183

Chemguide Entropy


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Periodicity

Elements of Period 3

Na Mg Al Si P S Cl Ar

Structure giant metallic Giant

Covalent

Simple Molecular Atoms

Electrical

Conduction

good, delocalised (mobile) e-

increases Na Al as more e- delocalised

poor, electrons localised in covalent

bonds so not free to move through

structure

poor

Melting Point high, strong metallic

bonding, strong forces of

attraction throughout entire

structure increases Na Al as more e- delocalised

higher,

strong

covalent

bonding

between

atoms

low, strong covalent bonds

between atoms, but weak

intermolecular forces

between Molecules

very

low

Mpt/oC 98 651 660 1410 44 113 -101 -189

silicon has a similar structure to diamond each silicon has 4 covalent bonds

silicon is important in the electronics industry

sodium is used in streetlights and as a coolant in nuclear reactors

magnesium and aluminium are important structural materials often used in alloys

chlorine is a product of the salt-alkali industry and has numerous uses

argon is used in lighting and lasers

sulphur is a raw material for the manufacture of sulphuric acid

Reaction of the elements of Period 3 with water

Na vigorous reaction darts about the surface dissolves to form a very alkaline

solution

Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

Mg

very slow with cold water, but

more readily with steam

Mg(s) + H2O(l) MgO(s) + H2(g)

Cl chlorine reacts with water to produce a

mixture of two acids by disproportionation

(moist blue litmus is first turned red by

chlorine gas then bleached)

0 -1 +1

Cl2(g) + H2O(l) HCl(aq) + HOCl(aq)

Unit 2

revision

only


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Reactions of the Elements of Period 3 with Oxygen

reactivity (generally) decreases across the period

argon does not react at all

chlorine-oxygen compounds do exist but are not formed by direct combination

the metals become coated in an oxide layer hence appear dull

sodium must be stored under oil due to its reactivity

aluminium oxide forms a protective layer around aluminium protecting it from further corrosion

red phosphorus needs heating to react with oxygen but its allotrope, white phosphorus, is stored under water as when dry it burns spontaneously in oxygen

the oxide formed contains the element in its highest oxidation state except for sulphur

Na burns very vigorously with a

yellow

flame

4Na(s) + O2(g) 2Na2O(s)

Mg burns very vigorously with a white

flame white MgO ash

2Mg(s) + O2(g) 2MgO(s)

Al initially vigorous

2Al(s) + 3O2(g) 2Al2O3(s)

Si slow

Si(s) + O2(g) SiO2(s) silicon(IV) oxide

P vigorous white fumes of P4O10

P4(s) + 5O2(g) P4O10(s) phosphorus(V) oxide

S melts, blue flame forms SO2 a colourless choking acidic gas

S (s) + O2(g) SO2(g)

Summary Questions Page 191 1 3

A2 Chemistry AQA (Nelson Thornes) 186 - 188

Chemguide Period 3 elements


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Periodicity of Oxides

Period 3

Oxides

Na2O MgO Al2O3 SiO2 P4O10

SO3 (SO2)

Physical State

(298K)

Solid liquid

(gas)

Mpt/oC 1275 2852

higher size and charge of

cation

2072 lower since

different

structure

1610 580 ? 17 (-73)

Structure Giant Ionic Lattice (Al2O3 has a degree of covalency small ion

+ large charge)

Giant

Covalent similar to

diamond

Simple Molecular

Nature of

Oxide Basic Amphoteric Acidic

Reaction

of Oxide

Water Acid Alkali

Na Na2O + H2O 2NaOH oxide ion is hydrolysed

very soluble pH 14

acid + metal oxide salt + water

Mg MgO + H2O Mg(OH)2 a weakly alkaline solution

only slightly soluble pH 9

Al Al2O3 + 3H2O + 2OH-

2[Al(OH)4]-

Si SiO2 + 2OH- SiO32-

(silicate ion) + H2O

P P4O10 + 6H2O 4H3PO4 phosphoric(V) acid pH 0

P4O10 + 12OH- 4PO43-

(phosphate(V) ion) + 6H2O

S SO2 + H2O H2SO3 sulphuric(IV) acid pH 3

SO3 + H2O H2SO4 sulphuric(VI) acid pH 0

SO2 + 2OH- SO32-

(sulphite ion) + H2O

SO3 + 2OH- SO42-

(sulphate ion) + H2O

aluminium oxide acts as an abrasive (as corundum) due to the strong bonding

aluminium/magnesium oxides are used as a refractory material (furnace)

melting point silicon(IV) oxide > carbon dioxide (strong covalent bonding throughout the structure rather than simple molecular lattice held together by weak intermolecular forces)

you should also be able to write a 3 step equation for phosphoric acid and water

Summary Questions Page 193 1 3 Exam Style Questions Page 194 1 - 6

Page 249 7


Chemguide Period 3 oxides

Oxidation Numbers

inc. IMF


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oxidation is the LOSS of electrons, reduction is the gain of electrons OILRIG

oxidising agents (oxidants) accept electrons and are themselves reduced

reducing agents (reductants) donate electrons and are themselves oxidised

Oxidation State

this is a book keeping method of the effective control of electrons used in bonding

elements = 0

oxidation state of elements in simple ions = charge on ion

oxidation state of elements in polyatomic ions = charge on ion

oxidation state of elements of a compound = 0 the relatively more electronegative element is assigned the negative oxidation state

hydrogen = +1 (except in metal hydrides where it is -1)

oxygen = -2 (except in peroxide O22-) where it = -1)

group 1 metals = +1

group 2 metals = +2

fluorine = -1 (even with oxygen, which is +2 in OF2)

Aluminium = +3

metals are always positive in a compound or polyatomic ion

maximum possible oxidation state = group number (note not always possible for various reasons see later)

oxidation numbers and nomenclature e.g. cobalt(II) nitrate(V), phosphorus(V) oxide

take care not to mix up charge and oxidation numbers in polyatomic species e.g. a

sulphate(IV) ion does NOT have a 4- charge (the IV refers to the oxidation state of the

sulphur)

changes in oxidation numbers can be used to identify redox reactions in inorganic and organic reactions e.g. metal or halogen displacement reactions

we specifically refer to an element in a species (e.g. the iron in Fe2O3 is reduced)

assume that multiple instances of an element in a species have the same value (e.g. both carbons in ethene are -2)

OXIDATION oxidation number becomes relatively more positive

REDUCTION oxidation number becomes relatively more negative


Chemguide oxidation number


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Balancing Redox Equations (Acidic conditions)

protocol for constructing half equations

1 get the species correct (and number e.g. 2Cr3+)

2 balance the oxygen by adding H2O(l)s to the side with least Os 3 balance the hydrogens using H+(aq)s to the side with least Hs 4 balance the charge on each side by adding e- to relatively more positive side 5 add state symbols

try doing these important half equations (either way round):

(i) Fe3+(aq)/ Fe2+(aq) (ii) MnO4-(aq)/Mn2+(aq) (iii) Cr2O72-(aq)/Cr3+(aq) (iv) S4O62- (aq)/ S2O32-(aq) (tetrathionate/thiosulphate) (v) CO2(g)/C2O42-(aq) (ethanedioate ion)

now try combining these half equations:

(i) iodine and thiosulphate ions (ii) iron (II) ions and manganate(VII) ions (iii) iron(II) ions and dichromate(VI) ions (iv) manganate(VII) ions and ethanedioate ions (v) dichromate(VI) ions and ethanal

Summary Questions Page 199 1 2 (Q2 should have equilibria arrows) Exam Style Questions Page


Chemguide redox equation

HSW: This titration can be used to investigate

the % of copper in a coin your teacher will explain how if you ask

HSW: These titrations can be used to investigate

the % of iron in an iron tablet or iron in tea

(after some preparatory steps)

HSW: This titration can be used to standardise

solutions of potassium dichromate(VI) or

potassium manganate(VIII)

I have written

some answers. Try

to find them and do the bonus STAR question


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Redox equations in alkaline conditions

this is included to allow you to work out rather than rote learn redox reactions for some

organic (e.g. Fehlings and Tollens), transition metal reactions (e.g. chromium) in alkaline conditions and certain reactions associated with batteries

writing half equations for redox reactions under alkaline conditions is a little more tricky than under acidic conditions but here is a useful cheat

do it exactly as if it was in acidic conditions then cancel out the hydrogen ions by adding hydroxide ions equally to both sides

you must know and use the formula of the species containing the atom(s) to be oxidised/reduced as it exists in alkaline conditions

lets consider Fehlings (see Module 4) in which copper(II) ions are reduced to copper(I) which exists as copper(I) oxide and an aldehyde is oxidised to a carboxylic acid (why would

you not smell this?)

using the correct species for alkaline conditions but as if in an acidic solution you should arrive at:

RCHO(aq) + 2Cu2+(aq) + 2H2O(l) RCOOH(aq) + Cu2O(s) + 4H+(aq)

now the modification

add OH-(aq) to BOTH sides in order to cancel 2H+

subsequently cancel out H2Os to arrive at:

RCHO(aq) + 2Cu2+(aq) + 4OH-(aq) RCOOH(aq) + Cu2O(s) + 2H2O(l)

this method also works perfectly well if applied to half equations

try using it to derive the equations met in Periodicity of oxides reacting with alkali from the known reactant and product


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Electrode potentials and the Electrochemical Series

dynamic equilibrium exists between a metal and its ions in solution which results in a potential on the metal

Mn+(aq) + ne- M(s)

this potential depends upon the position of the equilibrium and cannot be measured absolutely

it will be different for a more or less reactive metal so there will be a potential difference, which is measurable, between the two pieces of metal (try chewing a piece of aluminium if

you have a metal filling youll get the idea)

the relatively more reactive a metal, the relatively more biased the above equilibria will be to the LHS, the relatively more negative the potential on that metal will be

you should know the practical arrangement of a single cell made up of two half cells (see figure 3 on page 200)

Zn2+(aq)|Zn(s) and Cu2+(aq)|Cu(s)

Note: the relatively more oxidised form is shown on the LHS by convention (in practice

both reactions can go either way depending upon what they are combined with)

the pieces of metal are called electrodes

(in this case the cathode is the +ve terminal unlike in electrolysis where its the ve terminal since the cathode is defined as the electrode where reduction takes place)

the wire allows electron flow between the two electrodes

the salt bridge provides mobile ions to complete the circuit potassium nitrate (or similar) used in the salt bridge as the ions are unlikely to partake

in electron exchange or precipitation

oxidation takes place at the negative electrode (zinc in this case)

reduction takes place at the positive electrode (copper in this case)

given that the copper is the +ve electrode (by + 1.10V in theory) electrons flow to it so the reactions that are taking place in the two

half cells are:

Cu2+(aq) + 2e- Cu(s) and Zn(s) Zn2+(aq) + 2e-

in effect the more reactive metal (better reducing agent) has displaced the less reactive metal from its solution


Chemguide redox equilibria, half cells


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Standard Hydrogen Electrode

there are a huge number of different half cells that can be constructed (some of which dont involve a metal in the redox reaction thus requiring an inert platinum electrode)

to have comparisons between them all would just be silly so we use an electrode potential equivalent of sea level to which all are then compared

this works exactly like in geography where relative heights of mountains and depths of ocean trenches are given allowing differences to be calculated

similarly, the standard electrode (reduction) potential Eo for each half cell is measured relative to the standard hydrogen electrode (SHE), which is defined as 0V under standard

conditions:

25oC

100Kpa

1.0 moldm-3 H+(aq) CARE must be strong acid, and 0.5 moldm-3 if H2SO4

a diagram of the standard hydrogen electrode (PRIMARY STANDARD) and how it is used is shown in figure 6 on page 201 which you must be able to draw

a platinum electrode is used as this will not undergo a redox reaction thus will not alter the potential of the hydrogen electrode

it is covered with finely divided platinum (termed platinized platinum) to increase surface area and increase rate

all solutions will be 1.0 moldm-3 and 25oC as other values will alter the relative position

of the equilibria and therefore the absolute and hence measured potential difference

in practice the Calomel electrode A SECONDARY STANDARD is used for convenience

and is initially calibrated against the SHE at a value of +0.27V

Hg2Cl2(aq) + 2e- 2Hg(l) + 2Cl-(aq)

the voltmeter used has a high resistance hence a low current is drawn so a true reading of the potential difference is made

if a large current is drawn the concentrations of the ions in solution will change thus upsetting the equilibria and consequentially changing the value being measured


Chemguide hydrogen electrode,

non-metal systems, calomel


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Electrochemical Series

the electrochemical (reactivity) series is related to standard electrode potentials.

standard electrode potentials can be used to predict the relative feasibility of a reaction

by convention the standard electrode potentials are always quoted for the reduction process

the more +ve its value the relatively more likely the forward reaction is going to occur

i.e. the stronger the oxidant (on the left)

the more -ve its value the relatively more likely the reverse reaction is going to occur i.e.

the stronger the reductant (on the right)

It thus follows that

any species on the LHS can potentially oxidise any species on the RHS with a relatively

more negative electrode potential

any species on the RHS can potentially reduce any species on the LHS with a relatively

more positive electrode potential

thus relatively more reactive metals will have a relatively more ve electrode potential

lets use a familiar idea to reinforce this a more reactive metal can displace a less reactive metal e.g. iron nails in copper(II) sulphate solution (not the other way around)

Cu2+(aq) + 2e- Cu(s) Eo = + 0.34V

Fe2+(aq) + 2e- Fe(s) Eo = - 0.44V

which of the above metals will dissolve in 0.5M sulphuric acid?

NOTES:

1. the number of electrons lost or gained is not a factor in their relative availability 2 other factors, e.g. kinetics, may prevent a feasible reaction from occurring i.e

electrode potentials do not indicate the rate of a reaction, just its feasibility

(for example where a solid is involved, or where two ions of the same charge are the

reactants)


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Representing Cells

convention for representing cells (with the relatively +ve half cell on the RHS):

R O O R

Zn(s)| Zn2+

(aq)||Cu2+

(aq)|Cu(s)

single line is boundary between phases

commas are used between same phase (e.g. Fe3+(aq),Fe2+(aq))

double line = salt bridge (sometimes shown as a single dashed lines in some sources)

the emf quoted is the right hand side relative to the left hand side

Eocell = EoRHS - EoLHS

the hydrogen half cell can be represented in two different ways

life is easier if you put it on the side that provides a positive value for the cell overall as it is written

Pt(s)| H2(g)|H+(aq)||other half cell other half cell||H+(aq)|H2(g)|Pt(s)

e.g. for Zn2+(aq)|Zn(s) the electrons flow from the zinc half cell to the hydrogen half cell (Zn is higher in the reactivity series than hydrogen):

Zn(s)| Zn2+(aq)||H+(aq)|H2(g)|Pt(s)

Eocell = EoRHS - EoLHS +0.76V 0.00V - Zn half cell

the standard electrode potentials of ions of the same element in different oxidation states can be measured e.g. Eo Fe3+/Fe2+ = + 0.77V

1 What will be used as the electrode? 2. Show the cell diagram 3. What concentrations of iron sulphate solutions will you mix together?

Summary Questions Page 203 1 2 Exam Style Questions Page 213 4

A2 Chemistry AQA (Nelson Thornes) 203

Chemguide Electrochemical

e-

zinc is 0.76V

i.e its ( - 0.76)


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Electrode Potentials and Reaction Direction

the feasibility of a reaction requires electrons to flow in the right direction the reverse reaction is NOT feasible under standard conditions

R O O R

the oxidising agent must have a relatively more positive electrode potential

the reducing agent must have a relatively more negative electrode potential

remember that the electrode potentials are quoted for unimolar concentrations under standard conditions and will be modified through changes in:

concentration pH temperature

the effect of changes can be determined by applying LCP to the feasibility of a reaction e.g. increasing the concentration of one of the reactants involved in the oxidising half will

increase its electrode potential thus feasibility i.e. the forward reaction is more favoured

similarly increasing the concentration of one of the reactants involved in the reducing half

will decrease its electrode potential thus the reverse reaction is favoured

our technician wishes to prepare chlorine gas based on the electrode potentials given below

and you need to suggest how this can be achieved

Cl2 + 2e- 2Cl- Eo = +1.36 V

MnO2 + 4H+ + 2e- Mn2+ + 2H2O Eo = +1.23 V

when you study the variable oxidation states of chromium later on you will note that alkaline conditions are used during oxidation once again this modifies the associated electrode potentials and hence reaction feasibility

consider the feasibility of coating a copper/nickel coin with zinc as a first step to making a gold coin based on standard electrode potentials

during the operation of a battery the conditions change such that it goes flat

Summary Questions Page 207 1 4 Exam Style Questions Page 212 1 3 Page 246 3


Chemguide Electrochemical, vanadium, reaction

feasibility

e-

Reducing agent provides

e- and is itself oxidised

Oxidising agent accepts e-

and is itself reduced


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21

e-

Electrochemical Cells

1800 Alessandro Volta invents a primitive battery called a voltaic pile consisting of alternative layers of silver (or copper)

and zinc discs separated by cardboard soaked in salt water

Non-rechargeable Cells (Primary Cells)

The electrochemical reaction is not reversible (cells can be used only once) as when discharging the

cell the chemicals are permanently changed.

Daniell Cell

a wet cell invented by John Daniell in 1836

the porous pot was used (in place of the salt bridge shown below) to allow the ions to migrate when the battery was

operating whilst preventing the solutions from mixing

without this barrier, even when no current was drawn, the copper ions would be reduced at the zinc anode thus

shortening the battery's life

negative ions will migrate in the same direction as the electrons around the circuit and positive ions visa-versa

it was widely used at telegraph stations in the 19th century, however, its portability was limited due to the liquid electrolytes that it contained

Zn(s)| Zn2+(aq)||Cu2+(aq)|Cu(s)

Eocell = EoRHS - EoLHS 0.34 - - 0.76 = +1.10V

Note that, unlike in

electrolysis, the cathode

is the positive terminal

in a battery. This is

because the cathode is

specified as the electrode

at which reduction takes

place.

THE CATHODE Reduction takes place as the

positive pole of the battery

(half-cell with the highest

electrode potential)

accepts electrons from the

external circuit

Copper metal deposits on the

cathode hence its mass

increases

The solution becomes

more dilute hence the blue

colour fades

THE ANODE Oxidation takes place as the

negative pole of the battery

(half-cell with the lowest

electrode potential)

releases electrons to the

external circuit

Zinc goes into aqueous solution

(hence the zinc plate loses

mass)

Forgetting the double minus is a common

source of error

You could write

lots of notes but

ideally you

should be able to

work it all out

from this

schematic

or

ZnSO4

Learn the Daniel cell, zinc-carbon battery, lead acid battery and

hydrogen fuel cell. Develop an understanding of the rest.


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STANDARD Zinc Carbon Dry Cell

1887 - Carl Gassner patented a dry cell variant of the 1866 Leclanch wet cell

the electrodes are zinc and carbon, with an acidic paste between them

zinc serves as both the anode and the container, allowing the battery to be

completely self-contained and in effect more

portable and practical than wet cells

it could be used in any position as well rather than on a flat surface and the risk of leaking

was greatly reduced

the cathode is a mixture of powdered (surface area contact) manganese dioxide and graphite surrounding a solid graphite rod

the electrolyte is a paste of ammonium chloride inside the zinc can

overall reaction in a STANDARD zinc-carbon cell is:

Zn(s) + 2MnO2(s) + 2NH4+(aq) Mn2O3(s) + Zn(NH3)22+(aq) + H2O(l)

whilst zinc-carbon batteries are inexpensive they have very low power density so are only useful in devices that draw very little current

the zinc container also becomes thinner when used as the zinc is oxidised

it also thins when not used as ammonium chloride is acidic and slowly reacts with the zinc

which may lead to leakage

hence the service/shelf life of the battery is relatively short

the terminals of the battery are made of tin plated steel or brass to prevent the exposure the zinc, not allowing it to corrode as quickly, thus adding to the total battery life

the seal usually is made of asphalt pitch, wax, or plastic to allow the cathode mix (when the battery gets warm) to expand without rupturing the casing

- ANODE: Eo = -0.8V

Zn(s) Zn2+(aq) + 2e-

+ CATHODE: Eo = +0.7V

2NH4+(aq) + 2e- H2(g) + 2NH3(aq)

2NH3(aq) + Zn2+(aq) [Zn(NH3)2]2+(aq) prevents ammonia leaking (what shape is the complex ion)

+4 +3

2MnO2(s) + H2(g) Mn2O3(s) + H2O(l)

prevents a pressure build up from H2(g)

NH4+(aq) + H2O(l) H3O+(aq) + NH3(aq)

Note that overall

its the manganese that is reduced

This means

the casing

gets thinner

in use


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HEAVY DUTY Zinc Chloride Cell

zinc chloride cells are an improvement on the original zinc-carbon cell giving a longer life (about 50% due to a variation in

the chemical mix and greater mass)

electrolyte ZnCl2 paste (cf NH4Cl in standard Zn-C)

overall reaction in a HEAVY DUTY zinc-chloride cell is:

Zn(s) + 2MnO2(s) + H2O(l) Mn2O3(s) + Zn2+(aq) + 2OH-(aq)

these were originally marketed around 50 years ago as "Heavy Duty batteries, but since this term is not standardised it is a misleading as they are much inferior to alkaline batteries

which have since been introduced and are around 300% better

Alkaline Cell Battery

sold under brand names such as Duracell Energizer

these are more expensive than, but last considerably longer than, ordinary zinc-carbon cells

the cathode is manganese(IV) oxide powder

the anode is zinc powder (more surface area for increased rate of reaction therefore increased electron flow to allow for heavy duty usage

the electrolyte is potassium hydroxide paste (hence alkaline)

overall reaction is:

Zn(s) + 2MnO2(s) Mn2O3(s) + ZnO(s)

no gases (which insulate the electrodes) are produced which is one reason why they dont suffer from a voltage drop as do zinc-carbon batteries when worked hard

NOTE: Button batteries consist of a range of different types of system e.g. silver oxide, alkaline and are a structural form, for a specific use (compact size/long life), rather than a

particular chemistry so are not described in this guide. At one time mercuric oxide was used

but this has ceased for obvious reasons.

+ CATHODE:

+4 +3

2MnO2(s) + H2O(l) + 2e- Mn2O3(s) + 2OH-(aq)

+ CATHODE:

+4 +3

2MnO2(s) + H2O(l) + 2e- Mn2O3(s) + 2OH-(aq)

- ANODE:

Zn(s) + 2OH- (aq) ZnO(s) + H2O(l) + 2e-

Regenerated

- ANODE:

Zn(s) Zn2+(aq) + 2e-


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Rechargeable Cells (Secondary Cells)

The electrochemical reaction is not reversible (cells can be used only once) as when discharging the

cell the chemicals are permanently changed

Lead-acid battery

1859 - Gaston Plant - the lead-acid cell: the first rechargeable battery

the cathode is a lead-antimony alloy grid coated with lead dioxide

the anode is a lead-antimony alloy grid coated in spongy lead (its porous to increase

surface area)

the electrolyte is ~ 6M sulfuric acid in which the plates are immersed

at the anode lead combines with sulphate ions to

create lead sulfate and release electrons

as the battery discharges, both plates build up lead

sulphate and water builds up in the acid thus

diluting it

the voltage is about 2 volts per cell, so by

combining six cells in series you get a 12-volt

battery

upon discharging the following reactions take

place:

upon charging the reactions are reversed and this takes place when the car is running (using

the alternator) so that lead sulphate does not build up hence giving a flat battery even when not in use, leakage of current takes place so there is a net usage of sulphuric acid

hence the need for it to be checked and topped up when the vehicle is serviced

given the emphasis on fuel economy the battery must not add too much weight to the vehicle

but must provide enough power to start the car even in cold weather

for most of the last century these batteries have become standard for starting cars as they can

produce short bursts of a high current for many decades and are easy and cheap to

manufacture

however they cannot be used to power a vehicle for longer journeys as they provide relatively

little energy per kilogram and suffer power loss as insulating lead sulphate builds up

why do you think a car battery could explode if it is overcharged

- ANODE: Eo = -0.36V

Pb (s) + SO42-(aq) PbSO4(s) + + 2e-


PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- PbSO4(s) + 2H2O(l)

Overall: Eo = 2.04V

PbO2(s) + 4H+(aq) + 2SO42-(aq) + Pb(s) 2PbSO4(s) + 2H2O(l)


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Nickel-Cadmium Cells Nicads

1950s - Nickel-Cadmium (NiCd) first appeared

these are much more portable than lead-acid batteries but are more expensive

they are manufactured in sizes/voltages to act as direct replacements for cheaper zinc-carbon batteries

the fact that they can be recharged make them economical in the longer term

the cathode is made from nickel(III) oxyhydroxide

the anode is made from cadmium

the electrolyte is potassium hydroxide

when discharging the reactions are (reversed when charging):

cadmium is toxic so there are environmental issues regarding disposal

Nickel Metal Hydride Cells

1986 - NiMH battery was patented as a bi-product from research on the storage of

hydrogen for use as an alternative energy

source in the 1970s

some metallic alloys were observed to form hydrides that could capture (and release)

Hydrogen in volumes up to nearly a thousand

times their own volume

compared to lead-acid and NiCd, NiMH batteries have a higher storage capacity

they are more expensive than lead- acid and NiCd, but they are considered better for the

environment (lead and cadmium are toxic) prices are falling but lithium ion batteries are starting to gain some of the market

however, a NiCd battery has a lower self-discharge rate i.e. they hold their charge better

as with NiCds they offer slightly under 1.5V so may not work with some devices designed to operate off the 1.5V of alkaline or zinc-carbon batteries

the cathode is made from nickel(III) oxyhydroxide (the same as NiCds)


Cd(s) + 2OH-(aq) Cd(OH)2(s) + 2e-


NiO(OH)(s) + H2O(l) + e- Ni(OH)2 (s) + OH-(aq)

Overall: Eo = 1.33V

2NiO(OH)(s) + 2H2O(l) + Cd(s) 2Ni(OH)2 (s) + Cd(OH)2(s)

Note some sources state Ni(OH)3

here and omit H2Os to balance

Comparison of the discharge

voltage of an alkaline battery (red)

and a NiMH battery (blue). The green line is the voltage at which

the battery is considered dead


LUND 3 March 2015

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the anode is made from an alloy of rare earth metals that can soak up hydrogen atoms with the general formula AB5 where A can be lanthanum and B nickel

the electrolyte is potassium hydroxide

Lithium ion Cells

compared with NiMH and NiCd batteries of the same sizes or weights, Lithium Ion batteries are designed to deliver the highest energy output

a single cell voltage is 3.7V, 3 times that of NiMH batteries so a simpler battery configuration and better space utilization is achievable in devices such as cameras

they are relatively expensive as a computer chip is required to control charging and discharging but do offer a high capacity (hence reducing mass/size)

Li Ion batteries contain no toxic heavy metals, such as mercury, cadmium or lead

the cathode is made from lithium cobalt oxide powder

the anode is lithium/graphite formulation (a lot of technological development was required to prevent lithium oxidising and costing the electrode with insulating lithium

oxide

1996, the lithium ion polymer battery was developed from the lithium ion battery

these batteries hold their electrolyte in a solid polymer composite which cant leak

the electrodes and separators are laminated to each other with the whole devices encased in a flexible wrapping instead of a rigid metal casing, which means such batteries can be

specifically shaped to fit a particular device

Fuel Cells

The Hydrogen Economy

fossil fuels are at present the most economical way to power transportation

however, price rises commensurate with supply and demand, plus pollution issues such as the greenhouse gas CO2 and acidic nitrous oxides (from atmospheric N2 + O2 in a hot engine) etc

are driving the need for an alternative

an alternative fuel is hydrogen which if combusted does not give the pollution problems associated with hydrocarbons, the product is water


NiO(OH)(s) + H2O(l) + e- Ni(OH)2 (s) + OH-(aq)


AB5H(s) + OH-(aq) AB5(s) + H2O(l) + e-

Overall: Eo = 1.35V

NiO(OH)(s) + AB5H(s) Ni(OH)2 (s) + AB5


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hydrogen can be obtained from the electrolysis of sea water in the longer term but at present most hydrogen is still obtained from fossil fuels by steam methane reforming

this reacts steam with methane (natural gas) over a heated nickel catalyst to produce hydrogen and carbon monoxide

energy is obviously required to obtain the hydrogen so fuel cells are not the free energy from water that is often suggested

however, given that nuclear reactors cant be turned off, off peak generation could be one of the means of generating less expensive hydrogen along with wind, hydro, solar and tidal

whichever, there may still be some pollution associated with hydrogen production

Hydrogen Fuel Cell

in a battery the chemical energy is stored within the electrodes and the solution

in a fuel cell hydrogen (fuel) and air (oxygen) are fed into the cell in a similar way that petrol and air are fed into an internal combustion engine

the difference is that the chemicals are not combusted but react to produce electricity directly

this is more efficient than combusting the hydrogen (chemical heat kinetic electrical energy)

since a continuous supply of hydrogen is provided the voltage output remains constant

a typical fuel cell consists of two platinum electrodes

these also act as catalysts to assist the decomposition of hydrogen molecules

the electrodes are separated by a polymer electrolyte (proton exchange

membrane) through which hydrogen ions can

migrate whilst the gases are kept apart

a major issue is the storage and transportation of liquid hydrogen

research is currently being undertaken to develop hydrocarbon fuel cells so that car manufacturers can rely of normal fuel tanks

this is more complex as it requires preliminary reforming of the fuel within the vehicle

many other alternatives are currently under investigation

it is likely in the interim that fuel cell/battery/petrol hybrids will be employed to a greater extent

- ANODE: Eo = 0.0V

H2(g) 2H+(aq) + 2e- Zn(s) Zn2+(aq) + 2e-


4H+(aq) + O2(g) + 4e- 2H2O(l)

Overall: Eo = 1.2V

O2(g) + 2 H2(g) 2H2O(l)


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Overview

batteries reduce the need for expensive cabling and can provide power supplies to remote places

non-rechargeable batteries are cheap and can be manufactured in all sizes e.g. button batteries for watches, larger batteries for torches etc

however, they are thrown away afterwards along with energy and resources used in their manufacture

rechargeable batteries are more expensive initially but less resources are wasted in the longer term

they are vital in solar powered devices (and similar devices intent on storing power for future use)

the lead and nicads contain toxic chemicals thus there are disposal issues they must not go to landfill sites and must be recycled

neither are suitable for vehicles as they add too much mass and alternatives are being investigated

NiMH and lithium ion batteries are more environmentally friendly as they do not contain toxic heavy metals

sodium-sulphur batteries do offer a better power per kg output but have to operate at 300oC

another possibility is metal-air batteries (possible metals include aluminium and zinc)

a major issue is power density i.e. how much energy can be stored per kilogram of

battery, particularly where small size (e.g.

MP3 players) or small mass (transport) is

required

most batteries have performance that varies with temperature (either way)

fuel cells only produce water (spacecraft can use fuel cells to provide drinking water) and will eventually become the standard power source for vehicles which will reduce CO2

emissions IF the hydrogen can be produced cheaply without fossil fuels

they provide a more efficient means of converting chemical energy into electrical energy since it is direct rather than by a turbine

there are however problems associated with the transportation and storage of hydrogen and the means of refuelling the vehicle

you might be asked to calculate a voltage from electrode potentials but should be aware that the actual value will be unlikely to be this as it will not be operating under standard

conditions e.g variation in temperature

Exam Style Questions Page 246 3

A2 Chemistry AQA (Nelson Thornes) 200 201, 208 - 211

tbd


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Transition Metals

transition metal - d block elements that are able to form ions with a partially filled d sub shell

Sc is not a transition metal because its ion, Sc3+, is iso-electronic with Ar i.e. no d electrons

Zn is not a transition metal because its ion, Zn2+, has full d sub-shell

the electronic configuration for atoms and ions (remember to write 3d then 4s !) are written left to right in order of increasing energy

whilst the 4s subshell is initially of a lower energy than an unoccupied 3d, hence filled first,

adding electrons to the 3d pushes the 4s electrons away from the nucleus thus raising their

energy

dont forget that copper and chromium are not systematic

note that 4s electrons are always lost first when ions are formed and so first series transition metal ions never have any 4s electrons present so dont even show 4s (4s0 is incorrect)

if you are showing ions using the electrons in boxes nomenclature then note that paired d electrons are lost first (check Fe3+ and Co2+ for example) as mutual repulsion makes these

easier to remove

Physical Properties

typical metals i.e. malleable; ductile; good electrical and thermal conductors; all explained by the same ideas taught in Foundation Chemistry remind yourself of these.

melting point is higher than s block metals since there are more delocalised electrons holding structure together which also explains their greater mechanical strength

Chemical Properties

variable oxidation states

coloured ions (coloured rocks e.g. hmatite, malachite, typically include transition metal compounds)

catalysts

complex ions

Summary Questions Page 215 Q 1, 2


Chemguide Transition


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Complex Ions

a complex ion is a metal ion surrounded by ligands

ligands are molecules or ions which form dative (co-ordinate bonds) by donating electron pair(s) - lone pair donor - to a central metal ion lone pair acceptor

the coordination number is the number of coordinate bonds formed with the central metal ion NOT necessarily the number of ligands

ligands (electron pair donors) are Lewis bases

transition metal ions (electron pair acceptors) are Lewis acids

Unidentate ligands

form only one co-ordinate bond with the TM ion

NH3 ammine H2O aqua

OH- hydroxo Cl- chloro CN- cyano

SCN- thiocyanato

Pr- pavarotto (a rather large ligand not on the syllabus)

alphabetical order of name (prefixes i.e. di, tri etc ignored)

oxidation state of transition metal is given by roman numerals and this will only be the same as the

charge if all the ligands are neutral

name of metal is in Latin if complex ion is ve copper = cuprate lead = plumbate

iron = ferrate vanadium = vanadate

manganese = manganate chromium = chromate

zinc = zincate aluminium = aluminate

e.g. tetrachlorocuprate(II) ion [CuCl4]2- (ends in ate because it is anionic)

hexaaquacopper(II) ion [Cu(H2O)6]2+

tetraamminedicyanocobalt(III) ion [Co(NH3)4(CN)2]+

Neutral

Negative

number of each ligand

type of ligands central metal ion and its oxidation number

NOT

3+


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Bi-dentate ligands

two coordinate bonds per ligand

1,2-diaminoethane (en) benzene-1,2-diol

ethandioate (oxalate) ions (C2O42-)

benzene-1,2-dicarboxylate

two co-ordinate bonds per molecule leading to chelated (crab) complexes

this is also called chelation

the possibility of a bidentate ligand acting as a bridge between two separate metal ions exists

the replacement of unidentate with bidentate ligands is favoured by entropy since the total number of particles increases (see the section on entropy)

Multidentate ligands

more than two coordinate bonds per ligand

e.g. ethylenediamminetetraacetate ion - EDTA4-

EDTA complexes are very stable in effect a protective cage is formed around the transition metal ion thus isolating it from a biological system

the replacement of unidentate with multidentate ligands is favoured by entropy since the total number of particles increases (see the section on entropy)

1. antidote to Hg/Pb poisoning (traps metal ions) 2. Ca2+ trap in blood transfusions prevents clotting 3. removal of Ca2+ from hard water (e.g. in shampoo) 4. titimetric determination of metal ion concentration



Chemguide complex

Uses of

EDTA

Neutral

Negative


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Shapes of Complex Ions

octahedral e.g. with H2O, NH3, EDTA, en

whilst the coordination number is 6 there are 8 faces which determines the allocated name based on an octahedral unit

tetrahedral e.g. with Cl-

note that only 4 chloro ligands can fit around the central metal ion due to their relatively large size

hence the complex ion adopts a tetrahedral geometry

linear e.g. [Ag(NH3)2]+

(often in silver(1) Tollens reagent and copper(I) [Cu(NH3)2]+

complexes

square planar e.g. cis-platin [Pt(NH3)2Cl2]

(more on this later)

Extra info: geometrical and optical isomerism are possible

cis [Co(NH3)4Cl2] +

(aq) is violet

trans [Co(NH3)4Cl2] +

(aq) is green


Exam Style Questions Page 233 3


Chemguide Complex shape


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Colours of Complex Ions

light absorbed depends on the E of electronic transitions to the next vacant energy level

E = hh is Plancks constant

however, E(3d4s) in Ti3+ cannot account for lilac colour as E is too high (i.e. uv frequency)

in an isolated transition metal ion the d orbitals all have the same energy i.e. they are degenerate

however, ligands split the 3d energy level so that E is of a lower value corresponding to the energy of visible light

the reason for this is that the electrons donated by the ligand change the electronic environment to different extents for different d-orbitals in different geometrical positions i.e

the are all raised in energy but to differing degrees

white light incident upon a transition metal solution or solid will have certain wavelengths

absorbed in accordance with the value of E when exciting an electron thus removing this colour from the spectrum

colour observed is the complementary colour of light absorbed

hexaaquacopper(II) ions are blue as red is absorbed (see colour wheel)


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Sc3+ and Zn2+ are not coloured as there are no partially filled 3d sub-shells which is necessary for this to work i.e. to allow the promotion of d-electrons between d-ortbitals

colour depends on:

central metal ion this is a major factor - obviously

oxidation state e.g.

Fe(OH)2(s) Fe(OH)3(s) Cr2O72-(aq) [Cr(H2O)6]3+(aq) MnO4-(aq) Mn2+(aq)

[Co(NH3)6]2+(aq) [Co(NH3)6]3+(aq)

co-ordination number has a significant effect on d-d splitting hence colour change

varies size of E and the type of d-d splitting e.g.

[Co(H2O)6]2+(aq) [Co(Cl)4]2-(aq) [Cu(H2O)6]2+(aq) [Cu(Cl)4]2-(aq)

octahedral geometry yields 2 higher 3 lower

tetrahedral geometry yields 3 higher 2 lower

type of ligand stronger bonding causes greater d-d spitting hence shorter

wavelength absorbed

(spectrochemical series (Cl- < H2O < NH3 < en < CN-)

e.g.

[Cu(H2O)6]2+(aq) [Cu(NH3)4(H2O)2]2+(aq) [Co(H2O)6]2+(aq) [Co(NH3)6]2+(aq)

YOU MUST LEARN THESE COLOURS

[Cu(H2O)6]2+

(aq) Blue [Co(H2O)6]2+

(aq) Pink

[CuCl4]2-

(aq) Yellow [CoCl4]2-

(aq) Blue

Adding cHCl to [Cu(H2O)6]2+(aq) gives green! cobalt chloride paper is a test for water

[Cu(NH3)4(H2O)2]2+

(aq) Deep

Blue

[Co(NH3)6]2+

(aq) Yellow

NH3(aq) definitive test for Cu2+(aq) [Co(NH3)6]3+

(aq) Brown

[Fe(SCN)(H2O)5]2+

(aq) Blood

Red

sensitive test for the presence of Fe3+(aq)


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Colorimetry

you should re-familiarise yourself with the nature of light and why we see colour

two ways in which chemicals can interact with light result in:

absorption spectra (star light through a

a planetary atmosphere, chlorophyll)

emission spectra (e.g. flame tests, street lights)

absorption in the visible light varies according to the complex ion present, path length and

concentration

absorption of aqua complexes is relatively weak so colours are not very intense

certain complexing agents (e.g. EDTA) increase colour intensity to aid detection and determination

for example complexing [Fe(H2O)6]3+(aq) ions with colourless thiocyanate ions (SCN-) to produce the more deeply coloured [Fe(SCN)(H2O)5]2+(aq) complex ion which can be used to

detect low concentrations of iron in substances like tea by comparing absorbance against a

calibration curve of known concentrations.

in a colorimeter interchangeable filters are used to illuminate the sample with its complementary colour where absorption is greatest hence sensitivity optimised

it is also possible to determine the formula of a complex ion

as the complexing agent is added to separate batches of the transition metal sample the intensity of the colour will increase until there is no more transition metal ions for it to

combine with (the volume would be kept constant using water)

this allows us to determine the number of complex ions that combine with a transition metal where both concentrations are known (ideally the same value)

Summary Questions Page 222 1, 2

Exam Style Questions Page 233 3 (if not already done)


Chemguide Colour


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Variable Oxidation States

across the d block the effective nuclear charge increases

hence relative stability of 2+ oxidation state cf 3+ increases as the e- are better held

as the number of 4s/3d electrons increases from Ti to Mn so does the maximum oxidation state (Sc=3, Ti = 4, V = 5, Cr = 6, Mn = 7).

thereafter the maximum declines as effective nuclear charge increases suggesting that the 4s and unpaired 3d electrons only are involved

transition metals with charges > +3 cannot exist in aqueous solution where they exist as oxoanions instead e.g. MnO4-, Cr2O72-(aq), CrO42-(aq) with covalent bonding between the

oxygen and the transition metal (can you suggest their shape?)

this can be explained in two ways a lot of energy would be required to form a 4+ ion, and if it existed it would have a large charge density (thus be highly polarising) so would react

with water molecules and decompose them

+2 state tends to be reducing, as exemplified by Fe2+ in the manganate(VII) titration

some +2 ions are unstable in air due to aerial oxidation (where they are themselves reducing agents)

this can be pH dependent and occurs more readily in alkaline conditions e.g. keeping [Fe(H2O)6]2+(aq) in acidic solution helps it resist aerial oxidation to [Fe(H2O)6]3+(aq)

Redox Titrations

higher oxidation states (typically +4 and higher) are good oxidising agents

MnO4-(aq) and Cr2O72-(aq) are particularly good as oxidising agents in redox titrations

you will need to balance redox equations some revision may be necessary here

the titration procedure is pretty much the same as with acid-base titrations

acidic conditions are employed and you should be able to carry out a calculation to determine the minimum amount of sulphuric acid required

sulphuric acid is preferred (can you explain why each of the following: hydrochloric, nitric and ethanoic might not be suitable?)

Manganate(VII) shows a distinct colour change whilst dichromate(VI) require an indicator since both Cr2O72-(aq) and [Cr(H2O)6]3+ ions are coloured

the indicator used is sodium N-phenylamine-4-sulphonate which turns from colourless to purple at the end point

in case you wondered it works by changing colour at a particular electrode potential, in this

case + 0.84V


Exam Style Questions Page 233 1(some extra reading will be needed)


Chemguide Variable oxidation state, redox titration


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Chromium

Reducing Chromium (VI) to Chromium(III)

chromium in chromate(VI) can be reduced by reacting it with Zn in the presence of conc. HCl (this also releases H2 - a reducing atmosphere)

the feasibility can be demonstrated using electrode potentials

Cr2O72-

(aq) + 14H+

(aq) + 6e- 2Cr3+(aq) + 7H2O(l) Eo +1.33V

orange green

Cr3+(aq) + e- Cr2+(aq) -0.41V

green blue

Zn2+(aq) + 2e- Zn(s) -0.76V

the +2 state is readily oxidised back to the +3 state by air unless preserved in a reducing (e.g. H2) atmosphere

Ox State Chromium

+6 CrrO72- orange

CrO42- (shape?) yellow

+3 [Cr(H2O)6]3+ Green

+2 [Cr(H2O)6]2+ Blue

Oxidising Chromium (III) to Chromium(VI)

oxidation of transition metals tends to occur more readily in alkaline conditions

Iron(II) sulphate for example is kept in acidic conditions to prevent aerial oxidation

a plausible reason is that it is harder to remove electrons from the positively charged complex present in acidic solutions

Cr3+ can be oxidised to chromate(VI) by H2O2 in strongly alkaline conditions

initially further deprotonation in xs OH- produces deep green [Cr(OH)6]3-(aq)

[Cr(H2O)6]3+(aq) + 6OH(aq)- [Cr(OH)6]3-(aq) + 6H2O(l)

subsequent oxidation with H2O2 yields yellow chromate(VI) CrO42-(aq) ions (aka tetraoxochromate(VI)) ions)

The oxidation of cobalt

is covered later on


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writing half equations for redox reactions under alkaline conditions is a little more tricky than under acidic conditions but here is a useful cheat

do it exactly as if it was in acidic conditions then cancel out the hydrogen ions by adding hydroxide ions equally to both sides

i.e.

H2O2(aq) + 2H+(aq) + 2e- 2H2O(l)

cancelling 2H+, and subsequently H2O

H2O2(aq) + 2e- 2OH-(aq)

[Cr(OH)6]3-(aq) CrO42-(aq) + 2H2O(l) + 2H+(aq) + 3e-

cancelling 2H+, and subsequently H2O

[Cr(OH)6]3-(aq) + 2OH-(aq) CrO42-(aq) + 4H2O(l) +3e-

now balancing for electrons and combining:

2[Cr(OH)6]3-(aq) + 3H2O2(aq) 2CrO42-(aq) + 2OH-(aq) + 8H2O(l)

upon acidification orange dichromate(VI) Cr2O72-(aq) is formed this is an acid-base equilibrium NOT a redox check the oxidation state of chromium (you should be able to write the equation)

see if you can write half equations and then a full balanced equation for other oxidations carried out in alkaline conditions e.g.:

[Co(NH3)6]2+(aq) to [Co(NH3)6]3+(aq) by aerial oxygen

Co(OH)2(s) oxidised by H2O2 to Co(OH)3(s).

Fe(OH)2(s) to Fe(OH)3(s) by aerial oxygen


A2 Chemistry AQA (Nelson Thornes) 204 7, 226 228, 238

Chemguide Chromium


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Catalysis by Transition Metals

catalysts enable a different mechanism with a different activation energy (hence different rate) whilst being chemically unchanged at the end of a reaction

NOTE CATALYSTS CHANGE THE VALUE EA NOT THE SHAPE OF THE CURVE.

Heterogeneous Catalysis

heterogeneous catalysis the catalyst is in a different phase to the reactants

typically transition metals or their compounds are used e.g.

manufacture of ammonia Haber Process Fe

catalytic converters Pt and Rh

hardening fats (making margarine) Hydrogenation Ni

(adsorption onto the surface of the solid nickel catalyst weakens bonds) manufacture of nitric acid Ostwald Process Pt and Rh

manufacture of sulphuric acid Contact Process V2O5

adsorption occurs onto active sites and consequently:

weakens the bonds in the reactants hence lowers the activation energy

improves the stereochemistry for collisions by orienting molecules favourably

provides a localised relatively high concentration of reactants

adsorption must be strong enough to hold the reactant for long enough to promote a reaction but must not be too strong (e.g. as with tungsten) otherwise regeneration of active sites is too

slow as the product undergoes desorption from the surface


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in general the strength of adsorption decreases from left to right in the transition metals and the desorption occurs more readily in the same direction hence Fe (Haber process) Co and Ni

(Hydrogenation) are commonly used since they offer a compromise

Catalytic Converters

platinum and rhodium are coated onto a honeycomb ceramic material (minimises

costs whilst providing a large surface

area = increased rate) since adsorption

only occurs at the surface (expensive

metal underneath would be wasted)

the reactant gases form weak bonds with the surface of the catalyst (adsorption)

this weakens their bonds thus lowering the activation energy (additionally the catalyst also

helps promote more favourable molecular orientation)

this is followed by desorption in which the products depart

the catalyst selected provides bonding strong enough to hold the reactant gases on the surface whilst not preventing the products from leaving thus blocking an active site

CO and NO react to form CO2 and N2

NO also reacts with uncombusted hydrocarbons to produce CO2, H2O and N2

poisoning can occur if impurities contaminate the active sites e.g. sulphur dioxide and lead poisons catalytic converters hence unleaded low sulphur fuels must be used

in addition the finely coated Pt/Rh can be lost from the surface

this reduces efficiency and can result in an MOT failure and a large bill

Haber process

Pea sized Fe lumps are the catalyst - large surface area (enhanced by an aluminium oxide promoter) to

increase rate without requiring an even higher

temperature (energy cost plus unfavourable for

yield)

the iron catalyst does not effect the equilibrium position as both forward and backward reactions are

favoured equally

sulphur impurities (present in natural gas) can poison the iron so scrubbing is carried out to remove the sulphur compounds (carbon monoxide can also be a problem) however it

eventually has to be replaced

Write equations

for these


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Methanol Production

methanol is used as a chemical feedstock and as an additive to petrol

it can be manufactured by the reversible reaction between carbon monoxide and hydrogen in the presence of a copper catalyst or alternatively Cr2O3

the reactants (synthesis gas) are manufactured from the reaction of methane or propane with steam

Contact process

V2O5 used rather than faster Pt as lowers costs and less prone to poisoning

specific use of the variable oxidation states of transition metals is made

V2O5(s) + SO2(g) V2O4(s) + SO3(g)

V2O4(s) + 2

1 O2(g) V2O5(s)

Homogeneous Catalysis

homogeneous catalysis the catalyst is in the same phase to the reactants

in this case the reaction proceeds via an intermediate species and will typically have a two step reaction profile with two activation energies both less than that for the uncatalysed

reaction

same phase as reactants (e.g. all in solution):

acid catalysed esterification

enzymes in biological systems

chlorine free radicals (formed by the action of UV light on CFCs) and ozone (O3) depletion

Peroxodisulphate and iodide ions

redox reaction between peroxodisulphate (S2O82-) and iodide ions is slow as both are negatively charged

catalysed by iron(II) (or iron(III) either will do) note oppositely charged ions now react

Fe2+(aq) + S2O82-(aq) 2SO42-(aq) + Fe3+(aq)

2Fe3+(aq) + 2I-(aq) I2(aq) + Fe2+(aq)

this could be followed experimentally using a colorimeter

+5 +4

+4 +5 +

+2 +3

+3 +2

Reactants Products

Reactants Product


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Autocatalysis

some reactions can speed up rather than slow down relative to the initial rate

there can be several reasons for this:

an oxide layer on a surface is being removed before

the acid gets at the metal

the reaction might be exothermic

the product is itself a catalyst for the reaction this is autocatalysis

in all cases the reaction will eventually start to slow down as reactants are used up

e.g. manganate(VII) initially reacts slowly with ethanedioate ions (from oxalic acid)

2MnO4-(aq) + 5C2O42-(aq) + 16H+(aq) 2Mn2+(aq) + 10CO2(g) + 8H2O(l)

the Mn2+ ions produced autocatalyse the reaction, hence it actually speeds up once started

they change to Mn3+ initially but are changed back in the next step:

MnO4-(aq) + 4Mn2+(aq) + 8H+(aq) 5Mn3+(aq) + 4H2O(l)

2Mn3+(aq) + C2O42-(aq) 2Mn2+(aq) + 2CO2(g)

this could be followed experimentally using a colorimeter

Summary Questions Page 232 1 - 4

Exam Style Questions Page 233 1, 2, 4

Exam Style Questions Page 247 4, 8


Chemguide Heterogeneous catalysis

time

Conc.

Slow reduction in conc.

initially Faster reduction in conc. as

autocatalysis begins

Reaction slows down

as reagents run out


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Applications of Transition Metal Complexes

EDTA

this is an ethylenediamminetetraacetate ion or EDTA for short

EDTA complexes are very stable in effect a protective cage is formed around the transition metal ion thus isolating it from a biological system

the replacement of unidentate with multidentate ligands is favoured by entropy since the total number of particles increases (see the section on entropy)

1. antidote to Hg/Pb poisoning (traps metal ions) 2. Ca2+ trap in blood transfusions prevents clotting 3. removal of Ca2+ from hard water (e.g. in shampoo) 4. titimetric determination of metal ion concentration

Haemoglobin

haem forms four co-ordinate bonds (tetradentate) with Fe2+ (a porphyrin structure)

N in globin a protein forms a fifth to form haemoglobin

O2 or H2O form the sixth bond in oxyhaemoglobin or deoxyhaemoglobin

as oxygen is a poor ligand it is easily released in cells

lack of iron in the blood can cause anaemia as insufficient oxygen is transported resulting in tiredness and fatigue (or is that homework)

taking iron tablets which contain soluble iron(II) sulphate counteracts this

CO (which is a better ligand than oxygen) bonds with haemoglobin more strongly to form the relatively stable carboxyhaemoglobin thus reducing the bloods capacity to transport oxygen

cyanide ions act in a similar way

similar structures are found in a range of biologically important substances such as vitamin B12 (cobalt), and chlorophyll (magnesium)

Uses of

EDTA


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Cis-platin

cis-platin - [Pt(NH3)2Cl2] - is used in treating certain cancers (note the DNA complexing explanation on page 219 is wrong as its the chlorines that are displaced)

its full name is cis-diamminedichloroplatinum(II) in case you wondered

it forms DNA cross links via the platinum which damage the cancer cells

http://www.youtube.com/watch?v=Wq_up2uQRDo&feature=related

however, it does have side effects as it also effects normal cells e.g. renal toxicity, bone marrow suppression (loss of white blood cells increases the risk of other infections), fatigue

and hearing loss and can also induce nausea and vomiting.

testing renal function, blood and hearing is recommended before each cycle of therapy. so a cautious approach to dosage is necessary

http://www.cancerhelp.org.uk/about-cancer/treatment/cancer-drugs/cisplatin

geometrical isomerism possible the other form being trans-platin (which has no effect on cancer for stereochemical reasons

Tollens Reagent

diamminesilver(I) ion

[Ag(NH3)2]+(aq)

formed in ammonical silver nitrate (Tollens reagent)

used in silver mirror test for aldehydes and distinguish them from ketones

[Ag(NH3)2]+(aq) is reduced to Ag the silver mirror - and NH3 is displaced

see if you can write a balanced redox equation under alkaline conditions

complexing prevents the precipitation of Ag2O in alkaline conditions which would otherwise mask the test

[Ag(NH3)2]+(aq) is also formed when testing for silver chloride and silver bromide with the addition of ammonia following the silver nitrate test

the ligand displacement allows the precipitate to be re-solvated

A2 Chemistry AQA (Nelson Thornes) 69 70, 217 219, 241

Chemguide Cis-platin, haemoglobin, edta, Tollens


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Metal-aqua Ions

in aqueous solution tri-positive and di-positive TM ions form hexaaqua- complexes

ligands (electron pair donors) are Lewis bases and transition metal ions (electron pair acceptors) are Lewis acids

the presence of the ligand creates the familiar colour of transition metal solutions

this can also be locked into the crystalline structure again imparting colour (.xH2O)

e.g. white anhydrous copper(II) sulphate dissolves (very) exothermically in water to form blue hexaaquacopper(II) ions

[Cu(H2O)6]2+(aq) Blue

[Co(H2O)6]2+(aq) Pink

[Fe(H2O)6]2+(aq) Green

[Cr(H2O)6]3+(aq) Green

[Fe(H2O)6]3+(aq) Yellow

[Al(H2O)6]3+(aq) Colourless

Hydrolysis of Metal-aqua Ions

water ligands have increased O-H + bond polarity which promotes the abstraction of a hydrogen ion by another water molecule compared to that which takes place in the auto-

ionisation of water (see Kw)

[Fe(H2O)6]3+(aq) + H2O(l) [Fe(H2O)5OH]2+(aq) + H3O+(aq) Pale Lilac Orange oxonium ion

ACID BASE BASE ACID

this is hydrolysis (reaction with water) and makes the solution pH around 2 for a 1M solution

it would be wrong to assume that most populous species is the complex ion on the RHS as the equilibria still lies strongly to the LHS, however it does result in an increased hydrogen ion

concentration

further deprotonation very limited as water is a relatively weak base and the charge on the

complex ion is less positive and so the O-H + bond polarity is less pronounced.

the solution appears yellow as the orange colour is more intense than the pale lilac

relative acidity of M3+ cf M2+ reflects the relative polarising power of the central transition metal on the polarity of the O-H bond of the ligand

the chemistry of Al3+(aq) is similar to tri-positive transition metal ions


Page 243 1

How science works Page 235 Theories of acids


Chemguide Acidity of hexaaqua

Note the charge !

Actually they are lilac but the presence of a

small amount of orange [Fe(H2O)5OH]2+(aq) makes it appear yellow (see later)


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Reactions of Transition Metal-aqua ions with OH-(aq)

upon the addition of sodium hydroxide a metal hydroxide precipitation occurs

Cr(OH)3(s) green Cu(OH)2(s) pale blue

Fe(OH)3(s) orange Co(OH)2(s) blue goes brown on standing

Al(OH)3(s) white Fe(OH)2(s) green turns orange on standing

(some texts will show the water ligands as well)

deprotonation reactions occur with the addition of basic hydroxide ions to a greater extent than in aqueous solution and this eventually presents a neutral complex ion and thus

precipitation (hydroxide ions are a better base than water)

alternatively this reaction can be depicted as a reaction between a hydroxide ion and the oxonium ion (based on the equation on page 45) with a consequential shift in equilibria

either is acceptable as the outcome is essentially the same, but direct abstraction by OH-(aq) is easier to produce equations for.

[Cr(H2O)6]3+(aq) + OH-(aq) [Cr(H2O)5OH]2+(aq) + H2O(l)

[Cr(H2O)5OH]2+(aq) + OH-(aq) [Cr(H2O)4(OH)2]+(aq) + H2O(l)

[Cr(H2O)4(OH)2]+(aq) + OH-(aq) [Cr(H2O)3(OH)3](s) + H2O(l)

here there is no repulsion since there is no charge and so the complexes can hydrogen bond together producing a gelatinous precipitate.

all the hydroxide precipitates are solvated by the addition of acid which reverses the

equilibria (equations must be known acid + base salt + water).

precipitates insoluble in XS sodium hydroxide solution, but soluble in acid, are basic

Fe(OH)3(s) Cu(OH)2(s) Fe(OH)2(s) Co(OH)2(s)

XS NaOH(aq)

XS sodium hydroxide can cause the precipitate to re-dissolve for amphoteric hydroxides due to further deprotonation for:

Cr(OH)3(s) Al(OH)3(s)

[Cr(H2O)3(OH)3](s) + OH-(aq) [Cr(H2O)2(OH)4]-(aq) + H2O(l)

the charged particle created can now be solvated and if the sodium hydroxide solution is concentrated enough then the hexahydroxo- complex can eventually be formed.

[Cr(H2O)2(OH)4]-(aq) + OH-(aq) [Cr(H2O)(OH)5]2-(aq) + H2O(l)

[Cr(H2O)(OH)5]2-(aq) + OH-(aq) [Cr(OH)6]3-(aq) + H2O(l)

Any one of these ions would be

credited in the exam, but this one

is easiest to remember.


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Reactions of Transition Metal-aqua Ions in Solution with CO32-(aq)

carbonate ions react with oxonium ions to yield carbon dioxide gas

CO32-(aq) + 2H3O+(aq) CO2(g) + 3H2O(l)

the relatively high acidity of tri-positive transition metal results in their metal carbonates being unstable

[Fe(H2O)6]3+(aq) + H2O(l) [Fe(H2O)5OH]2+(aq) + H3O+(aq)

2H3O+(aq) + CO32-(aq) CO2(g) + 3H2O(l)

thus carbonate ions react with tri-positive transition metal ions to produce carbon dioxide gas HENCE FIZZING in addition to a hydroxide precipitate

the deprotonation equilibrium is shifted to the RHS as the carbonate ion removes the oxonium ion until the neutral triaquatrihydroxo- complex (the precipitate) is obtained

the overall equation should be known but can be derived on the basis of a shift in equilibrium as the carbonate ion reacts with oxonium ion (it is slightly harder to work it out starting with

the carbonate ion abstracting a hydrogen ion directly, but is also acceptable)

2[M(H2O)6]3+(aq) + 3CO32-(aq) 2[M(H2O)3(OH)3](s) + 3CO2(g) + 3H2O(l)

Cr(OH)3(s) green Fe(OH)3(s) orange Al(OH)3(s)

as previously it is soluble in acid since hydroxides are bases

it is not soluble in xs sodium carbonate solution as the concentration of hydroxide ions is relatively low so no further deprotonation occurs.

di-positive transition metal carbonates are stable as the oxonium ion concentration is relatively low

hence metal carbonate precipitates are produced and no CO2(g) is evolved

CoCO3(s) mauve CuCO3(s) blue-green FeCO3(s) green


Chemguide

Unit 5 Chemistry Notes

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