Medical ChemistryLecture 4 2007 (J.S.)

Oxidation-reduction reactions

Dissolution equilibriaPrecipitation reactions

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Oxidation-reduction reactions (redox reactions)

are common events in everyday life:

e.g. combustion of fossil fuels, corrosion of metals, dry cells and accumulators, metabolism of nutrients in human bodies, photosynthesis in green plants.

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Oxidation-reduction reactions are electron transfer reactions

The process may involve the complete transfer of electrons to formionic bonds or only a partial transfer or shift of electrons to formcovalent bonds.

Oxidation and reduction occur simultaneously in a chemicalreaction; one cannot take place without the other.

Ared – n e– Aox

Oxidationis the loss of electronsby a particle in a reaction,resulting in an increasein the oxidation number.

Box + n e– Bred

Reductionis the gain of electronsby a particle in a reactionthat results in a decreasein the oxidation number.

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Oxidation number (oxidation state of an element)

is the charge an element has in a simple ion it forms or the it would hypothetically have, if the shared electron pairs in covalent bonds are assigned to the more electronegative element sharing the pair of electrons.

Examples:

Oxidation number of sulfur (x) in sulfuric acid:

H2SO4

2 (+I ) 4 (–II ) 2 + (– 8) = – 6

– 6 + x = 0 + VI

Different oxidation numbers of nitrogen:

NH3–III N2

0 N2IO NIIO NIIIO2

– NVO3–

The algebraic sum of the oxidation numbers of elements in a molecularcompound equals zero and, in a polyatomic ion, it must equal the charge on the ion.

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Rules that help in assessing oxidation numbers:

– The oxidation number of any free element is zero, even when the atoms are combined with themselves (e.g. O2, P4, S8).

– No regard is paid to covalent bonds between atoms of the same species.

– An element may have more than one oxidation number, if it forms a variety of compounds.

– The oxidation number of hydrogen in a compound or an ion is + I except in ionic hydrides (– I).

– The oxidation number of oxygen in a compound or in an ion is –II except in peroxides (it takes on a – I).

– Metals generally have only positive oxidation numbers in compounds.

– The oxidation number of alkali metals equals always + I, of alkaline earth metals always + II.

– Nonmetals have negative oxidation numbers when combined with metals, positive oxidation numbers when combined with more electronegative nonmetals.

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Terms in redox reactions

A quite general reaction:

Ared + Box Aox + Bred

+ n e–

– n e–

In this reaction

Ared is oxidized because it loses electrons; it is a reductant (reducing agent) because it acts as donor of electrons and causes another species to be reduced.

Box is reduced because it gains electrons; it is an oxidant (oxidizing agent) because it acts as acceptor of electrons and causes another species to be oxidized.

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Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

+ 2 e–

– 2 e–

In the reaction

Fe is oxidized, it acts as reductant of Cu2+,

Cu2+ ion is reduced, it acts as oxidant of Fe.

Every redox reaction can be formally separated into two partscalled half-reactions (half-equations, half-cells) that representeither oxidation only or reduction only; they do not occur withoutthe other half-reaction taking place at the same time:

Fe Fe2+ + 2 e– oxidation

reduction Cu2+ + 2 e– Cu

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Pairs of the oxidized and reduced species Aox / Ared and Box / Bred that appear in half-equations are called

redox pairs (or redox couples).

Components of a particular redox pair can differ not only in the number of electrons but also in the number of hydrogen, oxygen, as well as other atoms.

Fe3+ / Fe2+

O2 / 2H2O MnO4

– / MnO2

Cr2O72– / Cr3+

aldehyde / alcohol pyruvate / lactate disulfide / 2 thioalcohols

quinone / diphenol

Examples of redox pairs:

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A redox pair is a couple of particles, which differ each from other in the oxidation number of one or more atoms of the same element (mostly also in the number of electrons). One component of a redox pair is more oxidized and can give the second one by reduction (in a "half-reaction" of a particular redox reaction).

Don't confuse redox pairs and conjugate pairs !

A conjugate pair is a couple that consists of an acid and a base(that differ just only in hydrogen ion H+ ).

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Redox reactions then may be easily identified not only through changes of charges on ions, but according to other symptoms:

Oxidation and reduction are defined generally in terms of changes in oxidation numbers. In some redox reactions, actual electron loss and gain to form ions need not to occur.

Hydrogenation and dehydrogenation are redox reactions, the products of which contain more or less hydrogen atoms

(as well as less or more multiple covalent bonds – the terms saturation or desaturation are also used).

Oxygenation and deoxygenation are redox reactions, theproducts of which contain more or less oxygen atoms.

A special type of oxygenation is hydroxylation.

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Cu2+ + Fe Cu + Fe2+ reduction of cupric ion to copper

Zn + Cu2+ Zn2+ + Cu oxidation of zinc

CO2 CO + ½O2 reduction of carbon dioxide by deoxygenation

C(s) + O2 CO2 oxidation (combustion) of carbon

Different types of redox reactions – examples:

– Loss and gain of electrons

– Oxygenation and deoxygenation

– Dehydrogenation and hydrogenation

CH3CH2-OH CH3CH=O dehydrogenation of ethanol to acetaldehyde– 2H

+ 2HCH3–C–COOH

O

CH3–CH–COOH

OH

hydrogenation (reduction)of pyruvate to lactate

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Do not confuse the terms hydrogenation and hydratation,

dehydrogenation and dehydratation !

Hydratation and dehydratation are not redox reactions; there is no change in the sum of the both carbon oxidation numbers (one of them is oxidized and another is reduced in addition or elimination of water).

CH CH–I –I –II

CH CH2

OH

0+H2O

–H2O

In organic chemistry, hydrogenated products are sometimes named by adding theprefix dihydro– to the name of a original compound, anddehydrogenated products by adding the prefix dehydro– to the name of a original compound.

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Well-known strong oxidants and reductants – examples:

Oxidizing agents – H2O2, KMnO4, K2Cr2O7, Cl2, I2

Reducing agents – H2, C, Fe, Zn, SnCl2

Oxidants and reductants differ in their ability to react with other agents considerably.

The strength of oxidants and reductants (their tendency togain or lose electrons) is expressed for particular redox pairsby standard electrode potentials E0.

Standard electrode potential E0 is the potential for an electrochemical half-cell (both oxidized and reduced form at c = 1 mol/l)

established relative to the potential of 0.000 V for the standard hydrogen electrode (H+/H couple under standard state conditions).

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Standard state of a half-cell: – both oxidized and reduced form of a redox pair at c = 1 mol/l – specified temperature, usually 25 °C – atmospheric pressure 101.3 kPa is important only when there is a gaseous component of the redox pair

H2

H+

[H+] = 1 mol / lpH2 = 101.3 kPa

E0(H+/H) = 0.000 V (25 °C)

Standard hydrogen electrode– reference electrode

electrode -an inert metal

Aox

Ared

[Aox] = [Ared] = 1 mol/l

Half-cell to be measuredin the standard state

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The equilibrium electromotive force Ecell (the potential of the galvanic cell) that is the potential difference between the twohalf-cells is measured:

H+

H2

Aox

Ared

salt bridge

millivoltmeter withhigh inner resistance

Any other reference electrode (which is stable and the potential known)may be used for measurement of electrode potentials, e.g. silver chloride or saturated calomel electrode (E0 = + 0.246 V).

E0cell = ΔE0 = E0

X – E0ref

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Half-reaction E° (V)

Na+ + e− Na

Zn2+ + 2 e− Zn

Fe2+ + 2 e– Fe

2 H+ + 2 e− H2

Cu2+ + 2 e− Cu

I2 + 2 e− 2 I −

Fe3+ + e− Fe2+

O2 + 4 H+ + 4 e− 2 H2O

Cr2O7− + 14 H+ + 6 e− 2 Cr3+ + 7 H2O

MnO4− + 8 H+ + 5 e− Mn2+ + 4 H2O

H2O2 + 2 H+ + 2 e− 2 H2O

− 2.71

− 0.76

– 0.44

0.00

0.34

0.54

0.76

1.23

1.33

1.51

1.77

Examples of standard electrode potentials (25 °C)

at pH 0(at pH 7.0 E°´= – 0.420)

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Redox pairs from the previous table:

The guiding principle:Under standard conditions, any oxidant will react with anyreductant with a lower, more negative E0 (i.e. situatedabove in the table).

Na+ / NaZn2+ / ZnFe2+ / FeH+ / ½ H2

Cu2+ / CuI2 / 2 I−

Fe3+ / Fe2+

O2 / 2 H2OCr2O7

2− / 2 Cr3+

MnO4− / Mn2+

H2O2 / H2O

strong oxidants

Oxidized forms areweak oxidants

Reduced forms arestrong reductants

weak reductants

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If the difference ΔE0 between both redox pairs is greater than 0.400 V, the reaction is irreversible (i.e. proceeds to completion) even under various non-standard concentrations of the reactants.

If the difference between both E0 is less than 0.400 V, then the reaction will reach equilibrium, the position of which depends on the initial concentrations of components of both redox pairs;the direction of such a reaction may be reversed.

Electrode potentials E under non-standard conditions

for a redox pair a Aox + n e– b Ared

Nernst equation

RTnF

lnE = E0 +[Aox]a

[Ared]b

[Aox] and [Ared ] relevant concentrations of reactants

R = 8.314 kPa K–1 mol–1

F = 96 500 C mol–1

n = number of moles of electrons transferred

E, E0 el. potentials in volts

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0.059n logE = E0 +

[Aox]a [Ared]b

After expressing R, T (298 K), and F in numbers and transposing natural logarithm into decadic (ln x = 2.3 log x), the equation will take the form

(in volts; t = 25 °C)

The electrode potential of half-cells at various concentrations of redox pair components can be calculated. On the contrary, the ratio of both redox pair components can be estimated from the measured values of electrode potentials.

Galvanic cells (electrochemical cells) are two half-cellsconnected by an external conductor.This arrangement allows discharging of the cell,the spontaneous cell reaction (the sum of the chemical half-reactions).

The electrons lost flow in the external circuit from the substance that is being oxidized in one of the half-cells to the substance that is being reduced in another half-cell till the cell reaches an equilibrium.

The free energy resulting from the spontaneous reaction is released as electrical energy and/or heat.

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Before a conductive connection of the half-cells, the electromotive force of the galvanic cell Ecell equals

Ecell = ΔE = EA – EB .

The expected cell reaction is

According to Nernst equation,

Ared + Box Aox + Bred

The cell potential ΔE is sometimes described as the "driving force" of the cell reaction. It is related to the amount of electrical work that a cell can perform.

the externalcircuit

After the external circuit is closed, the cell reaction is started. It goes on till the ΔE equals zero (the equilibrium state is reached).

Ecell = ΔE = ΔE0 +RTnF

ln[Aox]i [Bred]i

[Ared]i [Box]i

(in volts, 25 °C)

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A general rule suggests that the oxidized form of a redox pair with the more positive E is able to oxidize the reduced form of the redox pair with the less positive E.

that iodine will

Example:

Fe3+ + e– Fe2+ E20 = 0.75 V

I2 + 2 e– 2 I– E10 = 0.55 V

Conventionally, the more positive E in the cell is described as E2 .

Fe2+ + I2 Fe3+ + 2 I–Redox reaction:

Half-reactions:

The preferred reaction is then Fe2+ + I2 Fe3+ + 2 I–

The oxidation of iodide to elemental iodine will not be complete, the differenceΔE0 is less than 0.40 V.

Under which condition is iodine able to oxidize Fe2+ ?

ln[Fe3+]i [I– ]i

[Fe2+]i [½I2]i

ΔE = ΔE0 + RTnF

It follows from

oxidize Fe2+ to Fe3+ only if there is much higher initial concentration ofI2 and Fe2+ than of I– and Fe3+ (the ratio of the products [I]i [Fe2+]i and[I–]i [Fe3+]i must be greater than approx. 2500). Even so, the equilibriumwill be reached after oxidation of a very small amount of Fe2+.

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Relationship between the free energy change ΔG of redoxreactions and the cell potential ΔE

ΔG = ΔG0 + RT ln [A]ï [B]i

[C]i [D]i

The cell potential ΔE is a measure of whether a redox reaction isspontaneous. For the reaction Ared + Box → Aox + Bred it equals

The Gibbs free energy change ΔG is the quite general measure of reactionspontaneity and equals the free energy to do useful work. For the reactionA + B → C + D,

lnΔE = ΔE0 +

RTnF [Ared]ï [Box]i

[Aox]i [Bred]i

The electrical work available from a redox reaction –ΔG is equal to electrochemical potential ΔE times the electrical charge q (equal to nF) transferred in a redox reaction.

If a redox reaction starts at the standard state, – G0 = nF E0.

– G = nF E ( J mol–1 )

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Relationship beween ΔE0 and the equilibrium constant K

lnΔE = 0 = ΔE0 – RTnF [Ared]a

eq [Box]ceq

[Aox]beq [Bred]d

eq

The equilibrium constant Keq of a redox reaction

a Ared + c Box b Aox + d Bred is Keq = [Ared]aeq [Box]c

eq

[Aox]beq [Bred]d

eq

After the galvanic cell reaction reaches its equilibrium (the cell is discharged), ΔE equals zero ( E1 = E2 ).

ΔE0 = ln K and ln K = ΔE0

After expressing R, T (298 K), and F in numbers and transposing natural logarithm into decadic (ln x = 2.3 log x), the equation will take the form

log K = ΔE0 (25 °C)

RTnF

n0.059

RTnF

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Oxidation-reduction reactions in biological systems

Most biological redox reactions are catalyzed by enzymes.

Oxidative breakdown of nutrients rich in hydrogen supplies Gibbs free energy required to carry out various functions of living systems.

Some synthetic pathways, e.g. synthesis of fatty acids or cholesterol also include several redox reactions, but those are predominantly reductions (reductive syntheses).

Oxidation-reduction reactions serve to many other purposes, e.g. hydroxylations of numerous compounds foreign to the cells and dehydrogenation of alcohols.

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Photosyntheticautotrophs

Heterotrophsoxidize nutrients

CO2

H2O

O2

Nutrients rich in H

hν

Useful free energy in the form of ATP that drives endergonic reactions,

the remaining part of energy released as heat

The significance of biological oxidations in acquiring free energy

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Reduced cofactors NADH+H+ or FADH2 are reoxidized by giving over the pair of H atoms (called "a reducing equivalent") to the system of electron transporters of the terminal respiratory chain within the inner mitochondrial membrane.

The oxidation of nutrients is realized through several dehydrogenation steps.

Dehydrogenations are catalyzed by the enzymes dehydrogenases. The two atoms of hydrogen that are taken off from substrates are accepted by the oxidized forms of cofactors NAD+ or FAD.

+substrate

to be reducedoxidizedcofactor

+

2nd redox pair

1st redox pair

dehydrogenase oxidizedsubstrate

reducedcofactor

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NAD+ – the coenzyme of dehydrogenases (nicotinamide adenine dinucleotide)

It acts as an oxidant that takes off two atoms of hydrogen from the substrate.One atom plus one electron (hydride anion H–) is added to the para-position of the pyridinium ring, the remaining H+ binds to the enzyme.

Oxidized form NAD+

(aromatic ring, positive charge)

H

N

CONH2

H++H

Reduced form NADH + H+

(quinoid ring, no charge)

N

CONH2HH

+ H+

P–O–P

ribose

adenine

ribose

P–O–P

ribose

adenine

ribose

+2 H

-2 H

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Oxidized form of coenzyme FAD takes off two hydrogen atoms.

+2 H

-2 H

N

N

N

NH

O

O

H3C

CH2

FADOxidized form

H3C

CH2–O–P

H–C–OH

H–C–OH

H–C–OH

–O–P

ribose

adenine

N

N

N

NH

O

O

CH2

H–C–OH

H

H

FADH2

Reduced form

H3C

H3C

H–C–OH

H–C–OH

CH2–O–P –O–P

ribose

adenine

2 H

FAD – the coenzyme of dehydrogenases (flavin adenine dinucleotide)

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electrons

energy

4 H4 H+

2 O2–

CO2

Biological oxidations(dehydrogenations) Decarboxylations

Reduced coenzymes(reducing equivalents)

4e– 2 H2O

Nutrients rich in H

O2 + energy

Reduced coenzymes NADH+H+ FADH2

4 H+

NAD+ FAD

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Redox pairs in the terminal respiratory chain E°´ (V)

NAD+ + 2 H+ + 2 e− NADH + H+

FAD + 2 H+ + 2 e− FADH2

FMN + 2 H+ + 2 e− FMNH2

2 cytochrome b (Fe3+) + 2 e− 2 cytochrome b (Fe2+)

ubiquinone + 2 H+ + 2 e− ubiquinol

2 cytochrome c (Fe3+) + 2 e− 2 cytochrome c (Fe2+)

2 cytochrome a3 (Fe3+) + 2 e− 2 cytochrome a3 (Fe2+)

½ O2 + 2 H+ + 2 e− H2O

− 0.320a)

a)

0.030

0.100

0.235

0.385

0.816a) Flavoproteins exhibit variable values of E°´ (0.003 – 0.091 V) which depend on the protein part of the enzyme.

Electrode potentials in biological systems

are related to pH value 7.00 and temperature 30 °Cand then the symbols are the E 0' and E' instead of E0 and E.The standard potential of the hydrogen electrode at pH 7.00 is – 0.420 V,when compared to the hydrogen electrode at pH 0.00.

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H3C C

H

H

O

H

+ NAD H3C C

H

O+ NADH+H

Dehydrogenation of ethanol by alcohol dehydrogenase:

Additional examples of important oxidations-reductions

Four dehydrogenations in the citric acid cycle:

Isocitrate + NAD+ 2-Oxoglutarate + CO2 + NADH + H+

2-Oxoglutarate + NAD+ Succinyl-CoA + COě + NADH + H+

Succinate + FAD Fumarate + FADH2COOH

CH2

CH2

COOH

+ FAD

C

C

COOHH

HOOC H

+ FADH2

Malate + NAD+ Oxaloacetate + NADH + H

Reduction of pyruvate to lactate

Pyruvate + NADH + H+ Lactate + NAD+

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Dissolution equilibria

Precipitation reactions

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Let us recall the four various types of equilibria that may occur in aqueous electrolyte solutions:

1 Protolytic equilibria (acid-base equilibria)deal with the exchange of protons (hydronium ions)

between acids and bases

2 Oxidation-reduction equilibria

deal with the exchange of electrons in redox reactions

3 Dissolution equilibriaexpress relations of solids to ions and polar solvents insaturated solutions

4 Complex-forming equilibriaexist between donors and acceptors of valence shellelectron pairs

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3 Dissolution equilibria

Solubility of salts

When adding a salt tp water, afteraddition of certain amount a salt remainundissolved – the solution is saturated.

After certain time period, the equilibrium state will establish between the solid phase (undissolved salt)and the hydrated ions in thesaturated solution.

BnAm(s) n Bm+(aq) + m An–(aq)H2O

CaF2(s) Ca2+(aq) + 2 F–(aq)H2O

2+

2+

2+

CaF2(s)

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BnAm(s) n Bm+(aq) + m An–(aq)H2O

The equilibrium state of a saturated solution

CaF2(s) Ca2+(aq) + 2 F–(aq)H2O

or

is described by the equilibrium constant

[Bm+] [An–]KS(BnAm) = [Ca2+] [F–]2KS(CaF2) =

K = [Bm+]n [An–]m

[BnAm]

[Ca2+] [F–]2

[CaF2]K = or

Because the activity (then also "concentration") of any solid istaken as equal to unit ones and the amount of undissolved solidsalt does not influence the ion concentration in a saturated solution,

the simplified constant is called the solubility product KS :

e.g.,

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The solubility product indicates the maximal value of the product of ion concentrations (at a specified temperature). If the higher value than KS is reached, no matter in which way, the surplus of ions is separated out from the solution in the form of solid compound.

There are many precipitation reaction that are utilized in analytical chemistry and in syntheses of slightly soluble compounds.

The knowledge of calcium salts solubility helps in understanding the formation of renal stones, mineralization of bones and teeth, etc.

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Salt Ks (25 °C)

CaSO4

BaSO4

PbCl2

AgCl

CaF2

CaCO3

Ca(COO)2

CaHPO4

Ca3(PO4)2

Ca5(PO4)3F

1.2 10−6 1.4 10−10 1.6 10−5

1.8 10−10

2.7 10−11 3.8 10−9

1.0 10−9 2.3 10−7

2.8 10 −30

3.1 10−60

Solubility products of some slightly soluble salts

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A simplified survey of the solubility of ionic compounds is given in Medical Chemistry I, p. 80. Some relations between selected properties of metal ions and their valence electron configuration are described using the periodic table.

Insoluble hydroxides Amphoteric insoluble hydroxides

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Insoluble chlorides or other halides

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Insoluble sulfates

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The common ion effect

The common ion effect is the process, in which increasingthe concentration of one of the ions in an equilibrium resultsin a decrease of the concentration of the other ions.

An example:

To a saturated AgCl solution (KS = 1.7 10–10) solution of hydrochloricacid is added until the chloride concentration [Cl–] is 0.1 mol/l.

The concentration of silver ion [Ag+] in saturated solution equals 1.3 10–5.

KS = [Ag+] [Cl–] = 1.7 10–10 is the constant and remains the same nomatter how the concentration [Cl–] is changed.

Then KS = 1.7 10–10 = 0.1 x and x = [Ag+] = 1.7 10–9 after Cl– addition.

In this case, a 10 000-fold reduction of Ag+ ions in solution is shown.The extinct free Ag+ ions were precipitated as AgCl.

4 Complex-forming equilibria

Complex compounds originate in reactions, in which some cations oftransition metals bind donors of electron pairs (ligands with unshared electron pairs, e.g. ammonia, chloride or cyanide anions) through coordinate covalent bonds.

Under certain conditions, the complex particles dissociateto primary components. The stability of a complex particle, which is determined by the strength of the coordinate bond,may be expressed in the form of either the equilibrium dissociation constant of a complexor, more frequently, as the reciprocal of that dissociationconstant called the stability constant of a complex.