1 ENERGETICS
Simple treatment (including calculations) of the Born-Haber cycle for the halides of Group I and
(Lattice enthalpy will be regarded as the enthalpy of lattice breaking)
When you have finished this section you should be able to:
Explain and use the term ‘lattice enthalpy’ as a measure of ionic bond strength.
Construct Born-Haber cycles to calculate the lattice enthalpy of a simple ionic solid
e.g NaCl, MgCl2, using relevant energy terms (enthalpy changes of formation, ionisation energy, enthalpy of atomisation and electron affinity.
Explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy.
Lattice Enthalpy :
Bond enthalpies provide a measure of the strength of covalent bonds. Ionic bonding is an electrostatic attraction between oppositely charged ions.
The attraction acts in all directions, resulting in a giant ionic lattice containing many ions. For ionic compounds the corresponding enthalpy is the lattice enthalpy (also called the lattice energy).
Lattice enthalpy indicates the strength of the ionic bonds in an ionic lattice.
The standard molar lattice enthalpy is the energy required to convert one mole of a solid ionic compound into its constituent gaseous ions under standard conditions.
e.g. NaCl (s) Na+ (g) + Cl- (g)
Born-Haber cycle
This is an application of Hess’s Law and can be used to calculate lattice energies.
Lattice enthalpies cannot be determined directly by experiment and must be calculated indirectly using Hess’s Law and other enthalpy changes that can be found experimentally.
The energy cycle used to calculate a lattice enthalpy is the Born-Haber cycle.
The basis of a Born-Haber is the formation of an ionic lattice from its elements.
In general for an ionic compound a Born-Haber cycle can be written as:
M+ (g) + X- (g)
M+ (g) + X (g) ΔHe.a.
ΔHdiss.
M+ (g) + ½ X2 (g)
ΔHI.E. ΔHlatt.
M (g) + ½ X2 (g)
ΔHsub.
M (s) + ½ X2 (g)
ΔHFθ
M+X- (s)
According to Hess’s Law
ΔHlatt. = (-ΔHFθ) + ΔHsub. + ΔHI.E. + ΔHdiss. + ΔHe.a.
ΔHfθ = enthalpy of formation of MX (s)
ΔHsub. = enthalpy of sublimation of M (s)
ΔHI.E. = ionisation energy
ΔHdiss. = dissociation energy of X2 (g)
ΔHe.a.= electron affinity of X (g)
ΔHlatt. = lattice energy
N.B. The actual figures may be positive or negative and are simply substituted in the above equation.
Consider the reaction between sodium and chlorine to form sodium chloride.
Na (s) + ½ Cl2 (g) NaCl (s)
The reaction can be considered to occur by means of the following steps:
Vapourisation of sodium
Na (s) Na (g) ΔHsub.
The standard enthalpy of sublimation or vaporisation is the enthalpy change when one mole of sodium atoms are vaporised. This is an endothermic process and can be determined experimentally.
Ionisation of sodium
Na (g) Na+ (g) + e- ΔHI.E
The standard enthalpy of ionisation is the energy required to remove one mole of electrons from one mole of gaseous atoms. This is endothermic and can be determined by spectroscopy.
· Dissociation of chlorine molecules
Cl2 (g) 2Cl (g) ΔHdiss.
The standard bond dissociation enthalpy is the energy required to dissociate one mole of chlorine molecules into atoms (i.e. to break one mole of bonds). This is also endothermic and can be determined by spectroscopy.
Ionisation of chlorine atoms
Cl (g) + e- Cl- (g) ΔHe.a.
The electron affinity of chlorine is the energy released when one mole of gaseous chlorine atoms accepts one mole of electrons forming one mole of chloride ions.
Reaction between the ions
Na+ (g) + Cl- (g) NaCl (s) -ΔHlatt.
This is the reverse of the lattice energy. The standard lattice enthalpy is the energy absorbed when one mole of solid sodium chloride is separated into its gaseous ions. It has a positive value and cannot be determined experimentally.
A Born-Haber cycle can be drawn:
Na+ (g) + Cl- (g)
ΔHe.a.
Na+ (g) + Cl (g)
ΔHI.E.
Na (g) + Cl (g) ΔHlatt.
ΔHdiss.
Na (g) + ½ Cl2 (g)
ΔHsub. ΔHFθ
Na (s) + ½ Cl2 (g) NaCl (s)
Applying Hess’s Law
ΔHsub. + ΔHI.E. + ΔHdiss. + ΔHe.a. - ΔHlatt. – ΔHfθ = 0
Calculate the lattice enthalpy of sodium chloride given
ΔHfθ (NaCl) = -411 kJ mol-1
ΔHsub. (Na) = 108.3 kJ mol-1
ΔHI.E. (Na) = 500 kJ mol-1
ΔHdiss. (Cl) = 121 kJ mol-1
ΔHe.a. (Cl) = -364 kJ mol-1
Answer ΔHlatt. = +776 kJ mol-1
Exercise 1
Draw Born-Haber cycles for each of the following ionic compounds and calculate their lattice enthalpies.
(Note : ΔHat. of an element is the energy required to form one mole of gaseous atoms from the element.)
Solubility of ionic compounds is usually governed by its lattice energy. In general the higher the lattice energy the lower the solubility.
(See enthalpy of solution below)
A comparison of calculated and theoretical lattice energies gives an indication of the degree of covalent character in an ionic compound. The greater the difference between the two values the more covalent the compound.
The close agreement between the theoretical and experimental values for the alkali metal halides provides strong evidence that the simple ionic model of a lattice, composed of discrete spherical ions with an even charge distribution, is a very satisfactory one.
For the silver halides the theoretical values are about 15% less than the experimental values based on the Born-Haber cycle. This indicates that the simple ionic model is not very satisfactory.
When there is a large difference in electronegativity between the ions in a crystal , as in the case of the alkali metal halides then the ionic model is satisfactory. However as the difference in electronegativity gets smaller, as in the case of the silver halides, the bonding is stronger than the ionic model predicts.
The bonding in this case is not purely ionic but intermediate in character between ionic and covalent. The ionic bonds have been polarised (Fajans rules) giving some covalent character.
Exercise 2
The figures below give a list of lattice energies in kJ mol-1. Try to find as many patterns and trends in the figures as you can.
RbF -779 CaI2 -2038
BeF2 -3456 CaCl2- 2197
BaI2- 1841 MgCl2- 2489
MgBr2- 2416 KCl -710
CaBr2 -2125 NaF- 915
CsI- 607 LiF- 1029
KBr -671 MgI2 -2314
BaF2 -2289 LiBr- 804
CsBr -644 RbI- 624
LiI -753 SrBr2- 2046
BeI2 -2803 NaBr- 742
LiCl- 849 SrCl2- 2109
NaI -699 BeBr2- 2895
BeCl2- 2983 KF -813
CsCl -676 BaBr2- 1937
KI- 643 CaF2- 2583
MgF2 -2883 NaCl -776
RbCl- 685 SrF2 -2427
SrI2- 1954 RbBr -656
CsF -735 BaCl2- 2049
Trends in lattice enthalpy explained in terms of ionic radius and charge.
Consider the ionisation of an ionic solid MX.
MX (s) Mn+ (g) + Xn- (g)
The ease of separation of the ions and hence the lattice energy is determined by the size of the ions and their charge.
Effect of ionic size
As the ionic radius of both Mn+ and Xn- the lattice energy decreases. The attractive force between the ions decreases and they become easier to separate.
e.g. LiBr 804 kJ mol-1 BeCl2 2983 kJ mol-1
NaBr 742 MgCl2 2489
KBr 671 CaCl2 2197
RbBr 656 SrCl2 2109
CsBr 644 BaCl2 2049
As we descend both Groups I and II the lattice energies become less positive.
For any given metal the lattice energy also decreases in passing from the fluoride to the iodide.
e.g. NaF 915 kJ mol-1 SrF2 2427 kJ mol-1
NaCl 776 SrCl2 2109
NaBr 742 SrBr2 2046
NaI 699 SrI2 1954
This is due to an increase in ionic size from F- to I- which increases the internuclear distance. There is a corresponding decrease in attractive force and hence lattice energy.
When the internuclear distances are about equal, as for RbF and LiI for example, then the lattice energies are almost equal.
Effect of ionic charge
As the charge on Mn+ increases there is a greater attractive force between the ions and lattice energies increase. In addition, the decrease in size of Mn+ with increasing charge increases the attractive force between the ions and also increases the lattice energy.
The ionic radius of the Na+ and Ca2+ ions are very similar. However the lattice energy of CaCl2 is about 3 times that of NaCl.
NaCl 776 kJ mol-1 CaCl2 2197 kJ mol-1
This is due to the increased charge on the metal ion giving greater electrostatic attraction.
In general Group II halides have a lattice energy about three times that of the equivalent Group I halide.
Beryllium halides have considerable covalent character and the lattice energies are bigger than expected.
Exercise 3
What would be the effect on lattice energy of increasing the charge on Xn- ? (i.e. forming a Group VI compound rather than a Group VII compound).
Describe and explain the trends.
For comparable interionic distances the lattice energy would be bigger for X2- ions compared with X- ions. This is because X2- ions exert a stronger electrostatic field compared to X- ions due to the increased charge.
For comparable interionic distances the lattice energy is approximately four times higher when M2+ X2- ions are involved comared to M+ X-.
Exercise 4
Describe and explain the trends in hydration energy.
Sunday, April 19, 2009
As level Isomerism notes- work sheet- As chemistry work sheet
1. Which compound may exist as cis and trans isomers?
A CH3CH=CHCH3
B CH2=CHCl
C CH= CH2
D CH3 CH2
2. What is the total number of isomeric alkenes of formula C4H8?
A 2
B 3
C 4
D 5
3. Which one of the following molecules can exist as cis-trans isomers?
A (CH3)2C=CHCH3
B CH3CH2CH=CHCH3
C H2C=CHCH2CH2CH3
D CH3CH(CH3)CH=CH2
A CH3CH=CHCH3
B CH2=CHCl
C CH= CH2
D CH3 CH2
2. What is the total number of isomeric alkenes of formula C4H8?
A 2
B 3
C 4
D 5
3. Which one of the following molecules can exist as cis-trans isomers?
A (CH3)2C=CHCH3
B CH3CH2CH=CHCH3
C H2C=CHCH2CH2CH3
D CH3CH(CH3)CH=CH2
Isomerism in organic compounds
Isomerism in Organic Compounds
Structural isomerism for aliphatic compounds containing up to six carbon atoms, to include branched structures and position of the C=C double bonds in alkenes. (Cyclic compounds excluded).
Isomerism (greek isos, equal; meros, parts) occurs when there are two or more compounds (called isomers) with the same molecular formula but different arrangements of their atoms.
Isomers have different physical and chemical properties but the differences may be great or small depending on the type of isomerism.
There are two main classes of isomerism;
(i) structural isomerism
(ii) stereoisomerism
which are themselves sub-divided.
ISOMERISM
structural isomerism stereoisomerism
chain positional functional group geometric optical
isomerism isomerism isomerism isomerism isomerism
Structural isomerism
Structural isomers are molecules with the same molecular formula but with different structural arrangements of the atoms.
Chain isomers
This occurs where isomers have different arrangements of the carbon chain. There is only one alkane corresponding to each of the formulae CH4, C2H6 and C3H8.
For butane C4H10 two arrangements are possible.
n- butane
butane methylpropane
b.pt. 273K b.pt. 261K
As the number of carbon atoms increases the number of possible isomers increases rapidly.
Formula
Number of structural isomers
C5H12-3
C6H14-5
C7H16-9
Positional isomers
This occurs when isomers have the same carbon skeleton but the functional group is in different positions in the molecule.
propan-1-ol
propan-2-ol
Exercise 1
(a) Draw and name the five structural isomers of C6H14.
(b) Write the structural formulae of all the alcohols of molecular formula C4H10O. Name the isomers.
Stereoisomerism : cis-trans isomers for compounds containing one C=C bond, the energy barrier to rotation in these compounds.
Stereoisomerism
Stereoisomers have identical molecular formulae, and the atoms are linked together in the same order, but have different relative positions in space.
The two types of stereoisomerism are
(i) Geometric (or cis-trans ) isomerism
(ii) Optical isomerism [See Module 4.5]
· Geometric isomers
Geometric or cis-trans isomers exist because the π bond of the C=C bond prevents free rotation .
H H Cl H
C=C C=C
Cl Cl H Cl
cis-1,2-dichloroethane trans-1,2-dichloroethane
Cis-trans isomerism can also occur in inorganic complexes about a single bond with square planar or octahedral structures.
e.g. diaminedichloro platinum (II)
Exercise 2
1. Draw and name all the structural isomers, including geometric isomers, of C4H8.
Structural isomerism for aliphatic compounds containing up to six carbon atoms, to include branched structures and position of the C=C double bonds in alkenes. (Cyclic compounds excluded).
Isomerism (greek isos, equal; meros, parts) occurs when there are two or more compounds (called isomers) with the same molecular formula but different arrangements of their atoms.
Isomers have different physical and chemical properties but the differences may be great or small depending on the type of isomerism.
There are two main classes of isomerism;
(i) structural isomerism
(ii) stereoisomerism
which are themselves sub-divided.
ISOMERISM
structural isomerism stereoisomerism
chain positional functional group geometric optical
isomerism isomerism isomerism isomerism isomerism
Structural isomerism
Structural isomers are molecules with the same molecular formula but with different structural arrangements of the atoms.
Chain isomers
This occurs where isomers have different arrangements of the carbon chain. There is only one alkane corresponding to each of the formulae CH4, C2H6 and C3H8.
For butane C4H10 two arrangements are possible.
n- butane
butane methylpropane
b.pt. 273K b.pt. 261K
As the number of carbon atoms increases the number of possible isomers increases rapidly.
Formula
Number of structural isomers
C5H12-3
C6H14-5
C7H16-9
Positional isomers
This occurs when isomers have the same carbon skeleton but the functional group is in different positions in the molecule.
propan-1-ol
propan-2-ol
Exercise 1
(a) Draw and name the five structural isomers of C6H14.
(b) Write the structural formulae of all the alcohols of molecular formula C4H10O. Name the isomers.
Stereoisomerism : cis-trans isomers for compounds containing one C=C bond, the energy barrier to rotation in these compounds.
Stereoisomerism
Stereoisomers have identical molecular formulae, and the atoms are linked together in the same order, but have different relative positions in space.
The two types of stereoisomerism are
(i) Geometric (or cis-trans ) isomerism
(ii) Optical isomerism [See Module 4.5]
· Geometric isomers
Geometric or cis-trans isomers exist because the π bond of the C=C bond prevents free rotation .
H H Cl H
C=C C=C
Cl Cl H Cl
cis-1,2-dichloroethane trans-1,2-dichloroethane
Cis-trans isomerism can also occur in inorganic complexes about a single bond with square planar or octahedral structures.
e.g. diaminedichloro platinum (II)
Exercise 2
1. Draw and name all the structural isomers, including geometric isomers, of C4H8.
AS LEVEL -The Periodic Table
The organisation of elements in the Periodic Table according to their proton numbers and electronic structures. The terms group and period. The trends in the physical properties across the period sodium to argon limited to melting points, electrical conductivity, first ionisation energies and atomic radii.
Group VII (fluorine, chlorine, bromine and iodine)
Practical work restricted to chlorine, bromine and iodine and their compounds.
Trends within the group limited to colour, physical state, melting and boiling points, atomic and ionic radii, first ionisation energies, bond energies of halogen molecules, hydrogen halides and carbon-halogen bonds; electronegativies.
The halogens are a group of reactive non-metals, which are essentially similar to each other with only gradual changes as the atomic number increases.
Physical properties of the halogens
Property
Fluorine -F
Chlorine -Cl
Bromine -Br
Iodine -I
Colour
Pale yellow
Pale green
Red/brown
Black
Physical state
Gas
Gas
Liquid
Solid
They are all p-block elements with a simple molecular structure consisting of covalently bonded diatomic molecules, X2.
There are only weak Van der Waals forces between the molecules. The strength of the forces increases as the number of electrons (and Mr) in the molecule increases.
F2<> Br2 > I2 as the atoms get larger and the attraction of the nucleus for the shared electrons decreases (electronegativity decreases).
There is a slight tendency to metallic character with increasing atomic number. The halogens complete their octet by gaining one electron forming a halide ion, X- (see electron affinity values) or by sharing one electron.
Solubility in water and non-aqueous solvents eg hexane.
Solubility of the elements
All three elements are only slightly soluble in water because of the relatively strong hydrogen-bonding between the water molecules, which does not exist between the halogen molecules
i.e. solvent-solvent attractions > solute-solvent attractions > solute-solute
attractions.
Cl2> Br2>I2
solubility decreasing
They are soluble in non-polar organic solvents such as toluene and TCE.
(Why?)
Chemical trends: reactivity with hydrogen, sodium and phosphorus.
All the halogens are oxidising agents and combine readily with metals and non-metals.
Reaction of halogens with elements
Hydrogen
The halogens combine enthusiastically with hydrogen, the vigour of the reaction decreasing from fluorine to iodine.
H2 (g) + X2 (g) 2HX (g)
Fluorine reacts explosively even in the dark at –200 oC.
Chlorine reacts explosively in sunlight, or slowly in the dark below 200 oC.
Bromine reacts above 200 oC and at lower temperatures with a platinum catalyst.
Iodine reacts to form an equilibrium mixture
H2 (g) + I2 (g) ⇌ 2HI (g)
Metals
The halogens combine readily with most metals forming the metal halides.
The vigour of the reaction decreases from fluorine to iodine.
Group I and II halides are ionic.
2Na (s) + Cl2 (g) 2Na+Cl- (s)
Mg (s) + Cl2 (g) Mg2+2Cl- (s)
The halides of Group III are predominantly covalent.
2Al (s) + 3Cl2 (g) 2AlCl3 (s)
Non-metals
The elements react directly with many non-metals the oxidising power decreasing from fluorine to iodine.
The elements combine directly with phosphorus, the oxidation state of the product depending on the oxidising power of the halogen:
2P (s) + 5Cl2 (g) 2PCl5 (s)
2P (s) + 3Br2 (l) 2PBr3 (l)
Reactions of the elements illustrated by use of chlorine gas (or chlorine water), bromine water and aqueous iodine (in potassium iodide) with water, aqueous alkalis, other halides in solution and iron (II) and iron (III) ions as appropriate. Disproportionation.
Reaction of halogens with water
Fluorine and chlorine can oxidise water. Fluorine oxidises water to oxygen.
2F2 (g) + 2H2O (l) 4HF (aq) + O2 (g)
Chlorine reacts slowly with water forming hydrochloric acid and chloric(I) acid. This reaction involves disproportionation:- a change in which one particular molecule, atom or ion is simultaneously both oxidised and reduced.
reduction
Cl2 (g) + H2O (l) HCl (aq) + HClO (aq) (chlorine water)
o.n. 0 -1 +1
oxidation
Chlorine water contains chloric (I) acid HClO (aq), (hypochlorous acid). This is a weak acid which ionises to give the chlorate (I) ion ClO-, (hypochlorite ion). The hypochlorite ion is a powerful disinfectant and bleach.
Bromine disproportionates in a similar way but to a lesser extent.
Iodine has a very low solubility in water.
Reaction of halogens with aqueous sodium hydroxide.
Chlorine reacts faster with dilute sodium hydroxide than with water.
When chlorine is added to cold dilute alkali it disproportionates to chloride and chlorate(I).
(i)
reduction
Cl2 (g) + 2NaOH (aq) NaCl (aq) + NaOCl (aq) + H2O
o.n. 0 -1 +1
oxidation
( 2OH- + Cl2 Cl- + OCl- + H2O )
(ii) In hot concentrated alkali, if the solution is warmed to 70oC, the chlorate(I) disproportionates further to chlorate(V).
reduction
3NaOCl (aq) 2NaCl (aq) + NaClO3 (aq)
o.n. +1 -1 +5
oxidation
If chlorine is bubbled directly into hot conc. alkali then
(iii) reduction
3Cl2 (g) + 6NaOH(aq) 5NaCl (aq) + NaClO3 (aq)
o.n. 0 -1 +5
oxidation
( 6OH- + 3Cl2 5Cl- + ClO3- + 3H2O )
For bromine, both reactions (i) and (ii) are fast at 15oC.
For iodine, decomposition of IO- occurs rapidly at 0oC so it is difficult to prepare NaIO free from NaIO3.
NaClO is a mild antiseptic (Milton).
NaClO3 is a powerful weed killer.
Displacement reactions of the halogens
Since they are very electronegative, all the halogens are oxidising agents. As the group is descended their oxidising power decreases.
Therefore chlorine oxidises bromide ions to bromine and iodide ions to iodine.
These are displacement reactions.
Cl2 (g) + 2Br- (aq) Br2 (l)+ 2Cl- (aq)
(colourless) (yellow/orange)
Cl2 (g) + 2I- (aq) I2 (s) + 2Cl- (aq)
(colourless) (red/brown)
Bromine oxidises iodide to iodine
Br2 (g) + 2I- (aq) I2 (s) + 2Br- (aq)
Iodine does not oxidise any of the others.
Other oxidising reactions of the halogens
The trend in oxidising power is illustrated by the compounds formed by iron when it combines directly with the halogens.
Fluorine and chlorine form iron(III) fluoride and iron(III) chloride respectively.
2Fe (s) + 3F2 (g) 2FeF3 (s)
Bromine forms both iron(II) bromide and iron(III) bromide.
Iodine is too weak an oxidising agent and only forms iron (II) iodide.
Fe (s) + I2 (g) 2FeI2 (s)
Aqueous solutions of chlorine, bromine and iodine oxidise iron (II) to iron (III).
Cl2 (aq) + 2Fe2+ (aq) 2Cl- (aq) + 2Fe3+ (aq)
Iodine is so weak an oxidising agent that iron (III) ions oxidise iodide ions to iodine.
2Fe3+ (aq) + 2I- (aq) 2Fe2+ (aq) + I2 (s)
Thermal stability of hydrogen halides related to bond enthalpies. The relative strength of the acids, HF, HCI, HBr and HI.
Thermal stability of hydrogen halides
The thermal stability of the hydrogen halides decreases as the group is descended. This is in keeping with the trend in bond enthalpies
The size of the halogen atom increases from fluorine to iodine; therefore the bond length increases and the bond enthalpy decreases.
Hydrogen fluoride and hydrogen chloride are stable to heating. Hydrogen bromide decomposes on strong heating.
2HBr (g) H2 (g) + Br2 (g)
Hydrogen iodide decomposes on gentle heating. If a hot wire is dipped into hydrogen iodide gas violet clouds of iodine are produced.
2HI (g) H2 (g) + I2 (g)
Hydrogen halides as acids
The dry hydrogen halides are not acidic and do not affect litmus paper. The hydrogen halides dissolve readily in water forming acid solutions.
HX (g) + H2O (l) H3O+ (aq) + X- (aq)
The acid strength increases in the order
HF << HCl < HBr < HI
An aqueous solution of HF is weakly acidic. This is because the H-F bond is very strong and there is strong hydrogen bonding between the HF molecules.
The rest are all strong acids.
Ionic halides. The identification of halide ions in solution by use of silver ions followed by aqueous ammonia. The effect of light on silver halides. Presence of halide ions in sea water. The reaction of solid halides with concentrated sulphuric acid to illustrate the relative reducing ability of halides ions and hydrogen halide. The effects of fluoridation of public water supplies on dental health and an appreciation of the debate between public health policy and practice and the rights of the individual.
Reaction of the halide ions in solution, X-(aq)
Most metal halides are soluble except lead and silver halide. Therefore solutions of lead and silver ions are used to test for the presence of halide ions in solution.
Reagent
F- (aq)
Cl- (aq)
Br- (aq)
I- (aq)
Pb(NO3)2 (aq)
Pb2+(aq) + 2X-(aq) PbX2(s)
White precipitate of PbF2
White precipitate of PbCl2
Cream precipitate of PbBr2
Yellow precipitate of PbI2
AgNO3 (aq)
Ag+ (aq) + X- (aq) AgX (s)
No reaction AgF soluble in water
White precipitate AgCl
Cream precipitate AgBr
Yellow precipitate AgI
Solubility of silver halide in
(a) dil. NH3 (aq)
(b) conc. NH3
(c) dil.HNO3 (aq)
Exercise 2
Write an equation for the reaction of sodium chloride solution with
(a) lead nitrate solution and
(b) silver nitrate solution followed by the addition of ammonia.
Halide ions in sea water
Sea water contains over 3% of dissolved chlorides (mainly sodium, potassium, calcium and magnesium) and large solid deposits are found where inland seas have undergone evaporation. Bromine is also found in sea water as the bromides of sodium, potassium and magnesium. Iodine is present in sea water at a concentration of less than one part per million. Some seaweeds and sponges extract this iodine during growth and their ash contains about 0.5% by mass of iodine as iodides.
Reaction of the solid halides with conc. sulphuric acid, H2SO4
When concentrated sulphuric acid is added to a sodium halide the first product is fumes of the hydrogen halide HX, because each of these compounds is more volatile than sulphuric acid.
NaCl (s) + H2SO4 (1) HCl (g) + NaHSO4 (s)
NaBr (s) + H2SO4 (1) HBr(g) + NaHSO4 (s)
However conc. sulphuric acid is also a strong oxidising agent and will oxidise
HBr Br2 and HI I2, but not HF and HCl.
oxidised
2HBr(g) + H2SO4 Br2 + SO2 (g) + 2H2O
reduced
Similarly 2HI(g) + H2SO4 I2 + SO2(g) + 2H2O
Therefore conc. sulphuric acid cannot be used for the preparation of
HBr(g) and HI(g).
However conc. phosphoric(V) acid, H3PO4, can be used for the preparation since it is relatively non volatile and a poor oxidising agent.
NaBr (s) + H3PO4 (1) HBr (g) + NaH2PO4
NaI (s) + H3PO4 (1) HI (g) + NaH2PO4
Reagent
Fluoride
Chloride
Bromide
Iodide
Conc. H2SO4
X- + H2SO4 HX (g) + HSO4-
HF (g) colourless, pungent, corrosive gas
HCl (g)
formed
HBr (g) + a little Br2
A little HI (g) but mainly I2
Conc. H3PO4
HF (g)
HCl (g)
HBr (g)
HI (g)
Exercise 3
Make notes on the uses of the halogens and their compounds, with particular reference to the fluoridation of public water supplies and the effects on dental health.
KEY SKILLS
Group VII (fluorine, chlorine, bromine and iodine)
Practical work restricted to chlorine, bromine and iodine and their compounds.
Trends within the group limited to colour, physical state, melting and boiling points, atomic and ionic radii, first ionisation energies, bond energies of halogen molecules, hydrogen halides and carbon-halogen bonds; electronegativies.
The halogens are a group of reactive non-metals, which are essentially similar to each other with only gradual changes as the atomic number increases.
Physical properties of the halogens
Property
Fluorine -F
Chlorine -Cl
Bromine -Br
Iodine -I
Colour
Pale yellow
Pale green
Red/brown
Black
Physical state
Gas
Gas
Liquid
Solid
They are all p-block elements with a simple molecular structure consisting of covalently bonded diatomic molecules, X2.
There are only weak Van der Waals forces between the molecules. The strength of the forces increases as the number of electrons (and Mr) in the molecule increases.
F2<> Br2 > I2 as the atoms get larger and the attraction of the nucleus for the shared electrons decreases (electronegativity decreases).
There is a slight tendency to metallic character with increasing atomic number. The halogens complete their octet by gaining one electron forming a halide ion, X- (see electron affinity values) or by sharing one electron.
Solubility in water and non-aqueous solvents eg hexane.
Solubility of the elements
All three elements are only slightly soluble in water because of the relatively strong hydrogen-bonding between the water molecules, which does not exist between the halogen molecules
i.e. solvent-solvent attractions > solute-solvent attractions > solute-solute
attractions.
Cl2> Br2>I2
solubility decreasing
They are soluble in non-polar organic solvents such as toluene and TCE.
(Why?)
Chemical trends: reactivity with hydrogen, sodium and phosphorus.
All the halogens are oxidising agents and combine readily with metals and non-metals.
Reaction of halogens with elements
Hydrogen
The halogens combine enthusiastically with hydrogen, the vigour of the reaction decreasing from fluorine to iodine.
H2 (g) + X2 (g) 2HX (g)
Fluorine reacts explosively even in the dark at –200 oC.
Chlorine reacts explosively in sunlight, or slowly in the dark below 200 oC.
Bromine reacts above 200 oC and at lower temperatures with a platinum catalyst.
Iodine reacts to form an equilibrium mixture
H2 (g) + I2 (g) ⇌ 2HI (g)
Metals
The halogens combine readily with most metals forming the metal halides.
The vigour of the reaction decreases from fluorine to iodine.
Group I and II halides are ionic.
2Na (s) + Cl2 (g) 2Na+Cl- (s)
Mg (s) + Cl2 (g) Mg2+2Cl- (s)
The halides of Group III are predominantly covalent.
2Al (s) + 3Cl2 (g) 2AlCl3 (s)
Non-metals
The elements react directly with many non-metals the oxidising power decreasing from fluorine to iodine.
The elements combine directly with phosphorus, the oxidation state of the product depending on the oxidising power of the halogen:
2P (s) + 5Cl2 (g) 2PCl5 (s)
2P (s) + 3Br2 (l) 2PBr3 (l)
Reactions of the elements illustrated by use of chlorine gas (or chlorine water), bromine water and aqueous iodine (in potassium iodide) with water, aqueous alkalis, other halides in solution and iron (II) and iron (III) ions as appropriate. Disproportionation.
Reaction of halogens with water
Fluorine and chlorine can oxidise water. Fluorine oxidises water to oxygen.
2F2 (g) + 2H2O (l) 4HF (aq) + O2 (g)
Chlorine reacts slowly with water forming hydrochloric acid and chloric(I) acid. This reaction involves disproportionation:- a change in which one particular molecule, atom or ion is simultaneously both oxidised and reduced.
reduction
Cl2 (g) + H2O (l) HCl (aq) + HClO (aq) (chlorine water)
o.n. 0 -1 +1
oxidation
Chlorine water contains chloric (I) acid HClO (aq), (hypochlorous acid). This is a weak acid which ionises to give the chlorate (I) ion ClO-, (hypochlorite ion). The hypochlorite ion is a powerful disinfectant and bleach.
Bromine disproportionates in a similar way but to a lesser extent.
Iodine has a very low solubility in water.
Reaction of halogens with aqueous sodium hydroxide.
Chlorine reacts faster with dilute sodium hydroxide than with water.
When chlorine is added to cold dilute alkali it disproportionates to chloride and chlorate(I).
(i)
reduction
Cl2 (g) + 2NaOH (aq) NaCl (aq) + NaOCl (aq) + H2O
o.n. 0 -1 +1
oxidation
( 2OH- + Cl2 Cl- + OCl- + H2O )
(ii) In hot concentrated alkali, if the solution is warmed to 70oC, the chlorate(I) disproportionates further to chlorate(V).
reduction
3NaOCl (aq) 2NaCl (aq) + NaClO3 (aq)
o.n. +1 -1 +5
oxidation
If chlorine is bubbled directly into hot conc. alkali then
(iii) reduction
3Cl2 (g) + 6NaOH(aq) 5NaCl (aq) + NaClO3 (aq)
o.n. 0 -1 +5
oxidation
( 6OH- + 3Cl2 5Cl- + ClO3- + 3H2O )
For bromine, both reactions (i) and (ii) are fast at 15oC.
For iodine, decomposition of IO- occurs rapidly at 0oC so it is difficult to prepare NaIO free from NaIO3.
NaClO is a mild antiseptic (Milton).
NaClO3 is a powerful weed killer.
Displacement reactions of the halogens
Since they are very electronegative, all the halogens are oxidising agents. As the group is descended their oxidising power decreases.
Therefore chlorine oxidises bromide ions to bromine and iodide ions to iodine.
These are displacement reactions.
Cl2 (g) + 2Br- (aq) Br2 (l)+ 2Cl- (aq)
(colourless) (yellow/orange)
Cl2 (g) + 2I- (aq) I2 (s) + 2Cl- (aq)
(colourless) (red/brown)
Bromine oxidises iodide to iodine
Br2 (g) + 2I- (aq) I2 (s) + 2Br- (aq)
Iodine does not oxidise any of the others.
Other oxidising reactions of the halogens
The trend in oxidising power is illustrated by the compounds formed by iron when it combines directly with the halogens.
Fluorine and chlorine form iron(III) fluoride and iron(III) chloride respectively.
2Fe (s) + 3F2 (g) 2FeF3 (s)
Bromine forms both iron(II) bromide and iron(III) bromide.
Iodine is too weak an oxidising agent and only forms iron (II) iodide.
Fe (s) + I2 (g) 2FeI2 (s)
Aqueous solutions of chlorine, bromine and iodine oxidise iron (II) to iron (III).
Cl2 (aq) + 2Fe2+ (aq) 2Cl- (aq) + 2Fe3+ (aq)
Iodine is so weak an oxidising agent that iron (III) ions oxidise iodide ions to iodine.
2Fe3+ (aq) + 2I- (aq) 2Fe2+ (aq) + I2 (s)
Thermal stability of hydrogen halides related to bond enthalpies. The relative strength of the acids, HF, HCI, HBr and HI.
Thermal stability of hydrogen halides
The thermal stability of the hydrogen halides decreases as the group is descended. This is in keeping with the trend in bond enthalpies
The size of the halogen atom increases from fluorine to iodine; therefore the bond length increases and the bond enthalpy decreases.
Hydrogen fluoride and hydrogen chloride are stable to heating. Hydrogen bromide decomposes on strong heating.
2HBr (g) H2 (g) + Br2 (g)
Hydrogen iodide decomposes on gentle heating. If a hot wire is dipped into hydrogen iodide gas violet clouds of iodine are produced.
2HI (g) H2 (g) + I2 (g)
Hydrogen halides as acids
The dry hydrogen halides are not acidic and do not affect litmus paper. The hydrogen halides dissolve readily in water forming acid solutions.
HX (g) + H2O (l) H3O+ (aq) + X- (aq)
The acid strength increases in the order
HF << HCl < HBr < HI
An aqueous solution of HF is weakly acidic. This is because the H-F bond is very strong and there is strong hydrogen bonding between the HF molecules.
The rest are all strong acids.
Ionic halides. The identification of halide ions in solution by use of silver ions followed by aqueous ammonia. The effect of light on silver halides. Presence of halide ions in sea water. The reaction of solid halides with concentrated sulphuric acid to illustrate the relative reducing ability of halides ions and hydrogen halide. The effects of fluoridation of public water supplies on dental health and an appreciation of the debate between public health policy and practice and the rights of the individual.
Reaction of the halide ions in solution, X-(aq)
Most metal halides are soluble except lead and silver halide. Therefore solutions of lead and silver ions are used to test for the presence of halide ions in solution.
Reagent
F- (aq)
Cl- (aq)
Br- (aq)
I- (aq)
Pb(NO3)2 (aq)
Pb2+(aq) + 2X-(aq) PbX2(s)
White precipitate of PbF2
White precipitate of PbCl2
Cream precipitate of PbBr2
Yellow precipitate of PbI2
AgNO3 (aq)
Ag+ (aq) + X- (aq) AgX (s)
No reaction AgF soluble in water
White precipitate AgCl
Cream precipitate AgBr
Yellow precipitate AgI
Solubility of silver halide in
(a) dil. NH3 (aq)
(b) conc. NH3
(c) dil.HNO3 (aq)
Exercise 2
Write an equation for the reaction of sodium chloride solution with
(a) lead nitrate solution and
(b) silver nitrate solution followed by the addition of ammonia.
Halide ions in sea water
Sea water contains over 3% of dissolved chlorides (mainly sodium, potassium, calcium and magnesium) and large solid deposits are found where inland seas have undergone evaporation. Bromine is also found in sea water as the bromides of sodium, potassium and magnesium. Iodine is present in sea water at a concentration of less than one part per million. Some seaweeds and sponges extract this iodine during growth and their ash contains about 0.5% by mass of iodine as iodides.
Reaction of the solid halides with conc. sulphuric acid, H2SO4
When concentrated sulphuric acid is added to a sodium halide the first product is fumes of the hydrogen halide HX, because each of these compounds is more volatile than sulphuric acid.
NaCl (s) + H2SO4 (1) HCl (g) + NaHSO4 (s)
NaBr (s) + H2SO4 (1) HBr(g) + NaHSO4 (s)
However conc. sulphuric acid is also a strong oxidising agent and will oxidise
HBr Br2 and HI I2, but not HF and HCl.
oxidised
2HBr(g) + H2SO4 Br2 + SO2 (g) + 2H2O
reduced
Similarly 2HI(g) + H2SO4 I2 + SO2(g) + 2H2O
Therefore conc. sulphuric acid cannot be used for the preparation of
HBr(g) and HI(g).
However conc. phosphoric(V) acid, H3PO4, can be used for the preparation since it is relatively non volatile and a poor oxidising agent.
NaBr (s) + H3PO4 (1) HBr (g) + NaH2PO4
NaI (s) + H3PO4 (1) HI (g) + NaH2PO4
Reagent
Fluoride
Chloride
Bromide
Iodide
Conc. H2SO4
X- + H2SO4 HX (g) + HSO4-
HF (g) colourless, pungent, corrosive gas
HCl (g)
formed
HBr (g) + a little Br2
A little HI (g) but mainly I2
Conc. H3PO4
HF (g)
HCl (g)
HBr (g)
HI (g)
Exercise 3
Make notes on the uses of the halogens and their compounds, with particular reference to the fluoridation of public water supplies and the effects on dental health.
KEY SKILLS
AS LEVEL CHEMISTRY - PERIODIC TABLE NOTES
The Periodic Table
The organisation of elements in the Periodic Table according to their proton numbers and electronic structures. The terms group and period. The trends in the physical properties across the period sodium to argon limited to melting points, electrical conductivity, first ionisation energies and atomic radii.
Group VII (fluorine, chlorine, bromine and iodine)
Practical work restricted to chlorine, bromine and iodine and their compounds.
Trends within the group limited to colour, physical state, melting and boiling points, atomic and
ionic radii, first ionisation energies, bond energies of halogen molecules, hydrogen halides and carbon-halogen bonds; electronegativies.
GROUP VII Halogen
(Halogens : restricted to chlorine, bromine and iodine)
Valid deductions may be expected about other elements in the group.
Physical properties of halogens, limited to colour and physical state at room temperature.
The halogens are a group of reactive non-metals, which are essentially similar to each other with only gradual changes as the atomic number increases.
Physical properties of the halogens
Property
Fluorine
F
Chlorine
Cl
Bromine
Br
Iodine
I
Number of protons
(atomic number)
9
17
35
53
Outer electron configuration
2s22p5
3s23p5
3d104s24p5
4d105s25p5
Atomic radius/ nm
0.064
0.099
0.111
0.128
Ionic radius/ nm
0.133
0.181
0.196
0.219
Melting point / oC
-220
-101
-7
114
Boiling point / oC
-188
-34
58
183
Bond energy /
kJ mol-1
158
242
193
151
Electron affinity/ kJ mol-1
-361
-388
-365
-332
Standard electrode potential/ V
+2.87
+1.36
+1.09
+0.54
Electronegativity
4.00
2.85
2.75
2.20
Oxidation states
- 1
-1,1,3,5,7
-1,1,3,5,7
-1,1,3,5,7
Standard enthalpy of formation of NaX/ kJ mol-1
-573
-414
-361
-288
Standard lattice enthalpy of NaX/ kJ mol-1
902
771
733
684
They are all p-block elements with a simple molecular structure consisting of covalently bonded diatomic molecules, X2.
o o o o
o o o
o X o X o X X
o o o o
There are only weak Van der Waals forces between the molecules. The strength of the forces increases as the number of electrons (Mr) in the molecule increases.
F2<> Br2 > I2 as the atoms get larger and the attraction of the nucleus for the shared electrons decreases (electronegativity decreases).
There is a slight tendency to metallic character with increasing atomic number. The halogens complete their octet by gaining one electron forming a halide ion, X- (see electron affinity values) or by sharing one electron. Fluorine is restricted to an oxidation state of -1 but the remaining elements have empty d orbitals and can promote electrons to give oxidation states of +1, +3, +5 and +7.
They are all oxidising agents and combine readily with metals and hydrogen.
Chlorine is a greenish-yellow gas.
Bromine is a red-brown volatile liquid.
Iodine is a black shiny solid, which sublimes on heating to produce a purple vapour.
Solubility in water and non-aqueous solvents eg hexane.
Chemical trends: reactivity with hydrogen, sodium and phosphorus.
Reactions of the elements illustrated by use of chlorine gas (or chlorine water), bromine water and aqueous iodine (in potassium iodide) with water, aqueous alkalis, other halides in solution and iron (11) and iron (111) ions as appropriate.
Disproportionation.
Reaction of halides with elements
Metals
The halogens combine readily with most metals forming the metal halides.
The vigour of the reaction decreases from chlorine to iodine.
Group I and II halides are ionic.
2Na (s) + Cl2 (g) 2Na+Cl- (s)
Mg (s) + Cl2 (g) Mg2+2Cl- (s)
The halides of Group III are predominantly covalent.
2Al (s) + 3Cl2 (g) 2AlCl3 (s)
Non-metals
The elements react directly with many non-metals the oxidising power decreasing from chlorine to iodine.
The elements combine directly with phosphorus, the oxidation state of the product depending on the oxidising power of the halogen.
2P (s) + 5Cl2 (g) 2PCl5 (s)
2P (s) + 3Br2 (l) 2PBr3 (l)
Solubility of the elements
All three elements are only slightly soluble in water because of the relatively strong hydrogen-bonding between the water molecules, which does not exist between the halogen molecules
i.e. solvent-solvent attractions > solute-solvent attractions > solute-solute
attractions.
Cl2> Br2>I2
solubility decreasing
They are soluble in non-polar organic solvents such as toluene and TCE.
(Why?)
Chlorine reacts slowly with water forming hydrochloric acid and chloric(I) acid. This reaction involves disproportionation:- a change in which one particular molecule, atom or ion is simultaneously both oxidised and reduced.
reduction
Cl2 (g) + H2O (l) HCl (aq) + HClO (aq) (chlorine water)
o.n. 0 -1 +1
oxidation
Bromine and iodine disproportionate in a similar way but to a lesser extent.
Reaction of chlorine with aqueous sodium hydroxide.
Chlorine reacts faster with dilute sodium hydroxide than with water.
When chlorine is added to cold dilute alkali it disproportionates to chloride and chlorate(l).
(i) reduction
Cl2 (g) + 2NaOH (aq) NaCl (aq) + NaOCl (aq) + H2O
o.n. 0 -1 +1
oxidation
( 2OH- + Cl2 Cl- + OCl- + H2O )
(ii) In hot concentrated alkali, if the solution is warmed to 70oC, the chlorate(I) disproportionates further to chlorate(V).
reduction
3NaOCl (aq) 2NaCl (aq) + NaClO3 (aq)
o.n. +1 -1 +5
oxidation
If chlorine is bubbled directly into hot conc. alkali then
(iii) reduction
3Cl2 (g) + 6NaOH(aq) 5NaCl (aq) + NaClO3 (aq)
o.n. 0 -1 +5
oxidation
( 6OH- + 3Cl2 5Cl- + ClO3- + 3H2O )
For bromine, both reactions (i) and (ii) are fast at 15oC.
For iodine, decomposition of IO- occurs rapidly at 0oC so it is difficult to prepare NaIO free from NaIO3.
NaClO is a mild antiseptic (Milton).
NaClO3 powerful weed killer.
Thermal stability of hydrogen halides related to bond enthalpies. The relative strength of the acids, HF, HCI, HBr and HI.
Ionic halides. The identification of halide ions in solution by use of silver ions followed by aqueous ammonia. The effect of light on silver halides. Presence of halide ions in sea water. The reaction of solid halides with concentrated sulphuric acid to illustrate the relative reducing ability of halides ions and hydrogen halide. The effects of fluoridation of public water supplies on dental health and an appreciation of the debate between public health policy and practice and the rights of the individual.
Reaction of the halide ions in solution, X-(aq)
Most metal halides are soluble except lead and silver halide. Therefore solutions of lead and silver ions are used to test for the presence of halide ions in solution.
Reagent
F- (aq)
Cl- (aq)
Br- (aq)
I- (aq)
Pb(NO3)2 (aq)
Pb2+(aq) + 2X-(aq) PbX2(s)
White precipitate of PbF2
White precipitate of PbCl2
Cream precipitate of PbBr2
Yellow precipitate of PbI2
AgNO3 (aq)
Ag+ (aq) + X- (aq) AgX (s)
No reaction AgF soluble in water
White precipitate AgCl
Cream precipitate AgBr
Yellow precipitate AgI
Solubility of silver halide in
(a) dil. NH3 (aq)
(b) conc. NH3
(c) dil.HNO3 (aq)
soluble
soluble
insoluble
insoluble soluble
insoluble
insoluble insoluble
insoluble
Effect of sunlight
White ppt. turns purple/grey
Cream ppt. turns green/ yellow
No effect
Exercise 2
Write an equation for the reaction of sodium chloride solution with
(a) lead nitrate solution and
(b) silver nitrate solution followed by the addition of ammonia.
The organisation of elements in the Periodic Table according to their proton numbers and electronic structures. The terms group and period. The trends in the physical properties across the period sodium to argon limited to melting points, electrical conductivity, first ionisation energies and atomic radii.
Group VII (fluorine, chlorine, bromine and iodine)
Practical work restricted to chlorine, bromine and iodine and their compounds.
Trends within the group limited to colour, physical state, melting and boiling points, atomic and
ionic radii, first ionisation energies, bond energies of halogen molecules, hydrogen halides and carbon-halogen bonds; electronegativies.
GROUP VII Halogen
(Halogens : restricted to chlorine, bromine and iodine)
Valid deductions may be expected about other elements in the group.
Physical properties of halogens, limited to colour and physical state at room temperature.
The halogens are a group of reactive non-metals, which are essentially similar to each other with only gradual changes as the atomic number increases.
Physical properties of the halogens
Property
Fluorine
F
Chlorine
Cl
Bromine
Br
Iodine
I
Number of protons
(atomic number)
9
17
35
53
Outer electron configuration
2s22p5
3s23p5
3d104s24p5
4d105s25p5
Atomic radius/ nm
0.064
0.099
0.111
0.128
Ionic radius/ nm
0.133
0.181
0.196
0.219
Melting point / oC
-220
-101
-7
114
Boiling point / oC
-188
-34
58
183
Bond energy /
kJ mol-1
158
242
193
151
Electron affinity/ kJ mol-1
-361
-388
-365
-332
Standard electrode potential/ V
+2.87
+1.36
+1.09
+0.54
Electronegativity
4.00
2.85
2.75
2.20
Oxidation states
- 1
-1,1,3,5,7
-1,1,3,5,7
-1,1,3,5,7
Standard enthalpy of formation of NaX/ kJ mol-1
-573
-414
-361
-288
Standard lattice enthalpy of NaX/ kJ mol-1
902
771
733
684
They are all p-block elements with a simple molecular structure consisting of covalently bonded diatomic molecules, X2.
o o o o
o o o
o X o X o X X
o o o o
There are only weak Van der Waals forces between the molecules. The strength of the forces increases as the number of electrons (Mr) in the molecule increases.
F2<> Br2 > I2 as the atoms get larger and the attraction of the nucleus for the shared electrons decreases (electronegativity decreases).
There is a slight tendency to metallic character with increasing atomic number. The halogens complete their octet by gaining one electron forming a halide ion, X- (see electron affinity values) or by sharing one electron. Fluorine is restricted to an oxidation state of -1 but the remaining elements have empty d orbitals and can promote electrons to give oxidation states of +1, +3, +5 and +7.
They are all oxidising agents and combine readily with metals and hydrogen.
Chlorine is a greenish-yellow gas.
Bromine is a red-brown volatile liquid.
Iodine is a black shiny solid, which sublimes on heating to produce a purple vapour.
Solubility in water and non-aqueous solvents eg hexane.
Chemical trends: reactivity with hydrogen, sodium and phosphorus.
Reactions of the elements illustrated by use of chlorine gas (or chlorine water), bromine water and aqueous iodine (in potassium iodide) with water, aqueous alkalis, other halides in solution and iron (11) and iron (111) ions as appropriate.
Disproportionation.
Reaction of halides with elements
Metals
The halogens combine readily with most metals forming the metal halides.
The vigour of the reaction decreases from chlorine to iodine.
Group I and II halides are ionic.
2Na (s) + Cl2 (g) 2Na+Cl- (s)
Mg (s) + Cl2 (g) Mg2+2Cl- (s)
The halides of Group III are predominantly covalent.
2Al (s) + 3Cl2 (g) 2AlCl3 (s)
Non-metals
The elements react directly with many non-metals the oxidising power decreasing from chlorine to iodine.
The elements combine directly with phosphorus, the oxidation state of the product depending on the oxidising power of the halogen.
2P (s) + 5Cl2 (g) 2PCl5 (s)
2P (s) + 3Br2 (l) 2PBr3 (l)
Solubility of the elements
All three elements are only slightly soluble in water because of the relatively strong hydrogen-bonding between the water molecules, which does not exist between the halogen molecules
i.e. solvent-solvent attractions > solute-solvent attractions > solute-solute
attractions.
Cl2> Br2>I2
solubility decreasing
They are soluble in non-polar organic solvents such as toluene and TCE.
(Why?)
Chlorine reacts slowly with water forming hydrochloric acid and chloric(I) acid. This reaction involves disproportionation:- a change in which one particular molecule, atom or ion is simultaneously both oxidised and reduced.
reduction
Cl2 (g) + H2O (l) HCl (aq) + HClO (aq) (chlorine water)
o.n. 0 -1 +1
oxidation
Bromine and iodine disproportionate in a similar way but to a lesser extent.
Reaction of chlorine with aqueous sodium hydroxide.
Chlorine reacts faster with dilute sodium hydroxide than with water.
When chlorine is added to cold dilute alkali it disproportionates to chloride and chlorate(l).
(i) reduction
Cl2 (g) + 2NaOH (aq) NaCl (aq) + NaOCl (aq) + H2O
o.n. 0 -1 +1
oxidation
( 2OH- + Cl2 Cl- + OCl- + H2O )
(ii) In hot concentrated alkali, if the solution is warmed to 70oC, the chlorate(I) disproportionates further to chlorate(V).
reduction
3NaOCl (aq) 2NaCl (aq) + NaClO3 (aq)
o.n. +1 -1 +5
oxidation
If chlorine is bubbled directly into hot conc. alkali then
(iii) reduction
3Cl2 (g) + 6NaOH(aq) 5NaCl (aq) + NaClO3 (aq)
o.n. 0 -1 +5
oxidation
( 6OH- + 3Cl2 5Cl- + ClO3- + 3H2O )
For bromine, both reactions (i) and (ii) are fast at 15oC.
For iodine, decomposition of IO- occurs rapidly at 0oC so it is difficult to prepare NaIO free from NaIO3.
NaClO is a mild antiseptic (Milton).
NaClO3 powerful weed killer.
Thermal stability of hydrogen halides related to bond enthalpies. The relative strength of the acids, HF, HCI, HBr and HI.
Ionic halides. The identification of halide ions in solution by use of silver ions followed by aqueous ammonia. The effect of light on silver halides. Presence of halide ions in sea water. The reaction of solid halides with concentrated sulphuric acid to illustrate the relative reducing ability of halides ions and hydrogen halide. The effects of fluoridation of public water supplies on dental health and an appreciation of the debate between public health policy and practice and the rights of the individual.
Reaction of the halide ions in solution, X-(aq)
Most metal halides are soluble except lead and silver halide. Therefore solutions of lead and silver ions are used to test for the presence of halide ions in solution.
Reagent
F- (aq)
Cl- (aq)
Br- (aq)
I- (aq)
Pb(NO3)2 (aq)
Pb2+(aq) + 2X-(aq) PbX2(s)
White precipitate of PbF2
White precipitate of PbCl2
Cream precipitate of PbBr2
Yellow precipitate of PbI2
AgNO3 (aq)
Ag+ (aq) + X- (aq) AgX (s)
No reaction AgF soluble in water
White precipitate AgCl
Cream precipitate AgBr
Yellow precipitate AgI
Solubility of silver halide in
(a) dil. NH3 (aq)
(b) conc. NH3
(c) dil.HNO3 (aq)
soluble
soluble
insoluble
insoluble soluble
insoluble
insoluble insoluble
insoluble
Effect of sunlight
White ppt. turns purple/grey
Cream ppt. turns green/ yellow
No effect
Exercise 2
Write an equation for the reaction of sodium chloride solution with
(a) lead nitrate solution and
(b) silver nitrate solution followed by the addition of ammonia.
A LEVEL CHEMISTRY - PERIODIC TABLE
1.8 The Periodic Table
Section A
For each of the following questions only one of the lettered responses (A-D) is correct.
Select the correct response in each case and mark its code letter by connecting the dots as illustrated on the response sheet.
1.B Which one of the following equations does not represent a reaction of chlorine under suitable conditions?
A Cl2 + H2O HCl + HOCl
B 2Cl2 + CH4 CCl4 + 2H2
C 3Cl2 + 6NaOH 5NaCl + NaClO3 + 3H2O
D Cl2 + 2NaOH NaCl + NaOCl + H2O
2.A When the hydrogen halides dissolve in water acidic solutions are formed. Which has the highest pH (assuming all are equal concentrations)?
A Hydrogen fluoride
B Hydrogen chloride
C Hydrogen iodide
D Hydrogen bromide
3. In which one of the following changes has chlorine been oxidised?
A 3Cl2 + 2Fe 2FeCl3
B Cl2 + I2 2IC1
C Cl2 + 2KBr Br2 + 2KCl
D Cl2 + 2KOH H2O + KClO + KCl
4. In which one of the following sequences are the oxides of the elements classified as basic, amphoteric and acidic respectively?
A Na, Mg, Al
B Na, K, S
C K, Al, P
D S, P, Mg
5. Which one of the following is true of the halogens as the Group is descended from chlorine to iodine?
A The atomic radius decreases.
B The colour of the element lightens.
C The melting point of the element increases.
D The oxidising power of the element increases.
5. Which one of the following equations represents a reaction of chlorine?
A Cl2 + NaOH NaOCl + HCl
B Cl2 + Fe FeCl2
C Cl2 + H2O HCl + HOCl
D Cl2 + CH4 CH2Cl2 + H2
SECTION B
1. The group in the Periodic Table headed by fluorine is called the halogen group.
(a) Using your knowledge of the group predict the colour and physical state of astatine.
Black / solid
(b) When chlorine gas is bubbled through a solution of potassium iodide a redox reaction occurs.
(i) What is observed in this reaction?
The colour of the gas disappears / Solution turns colourless to dark brown
(ii) What is meant by a redox reaction?
A reaction in which one or more electrons / is~are transferred from one substance to another
(iii) Explain fully the changes undergone (if any) by the potassium and iodide ions.
Potassium ion : no change (1)/ iodide ion : loses electron to form iodine (1) [2]
(c)
(i) The presence of iodide ions in an aqueous solution of sodium iodide can be detected using silver nitrate solution and ammonia solution. Describe fully how you would carry out this procedure experimentally, stating any changes you observe. Write equations for any reactions which occur.
Add dilute nitric acid followed by silver nitrate solution to the iodide solution/ A pale yellow / precipitate / of silver iodide / forms. This is insoluble in ammonia solution. / AgNO3 + NaI = AgI + NaNO3 (1)
(ii) State what differences you would observe if the sodium iodide solution was replaced by sodium chloride solution. Write equations for any reactions which occur.
With sodium chloride a white precipitate of silver chloride forms. This is soluble in ammonia solution
2. Sodium bromide reacts with concentrated sulphuric acid at room temperature to give four products. Place ticks (√) in four of the boxes to show which substances are formed.
√ if the substance is formed
bromine
hydrogen
hydrogen bromide
hydrogen sulphide
sodium hydrogensulphate
sulphur
sulphur dioxide
3. The diagram below shows chlorine gas being passed through a dilute
solution of potassium iodide. The upper layer is a hydrocarbon solvent.
(a) (i) What is the most important safety precaution to take when carrying
out the experiment?
The colour of the gas disappears / Solution turns colourless to dark brown
(ii) Write the ionic equation for the reaction between chlorine and potassium iodide and explain the redox reactions taking place in terms of electron transfer.
With sodium chloride a white precipitate of silver chloride forms. This is soluble in ammonia solution
(b) When the aqueous layer is shaken with the hydrocarbon most of the iodine dissolves in the upper layer.
(i) What is the colour of iodine in potassium iodide solution?
(ii) What does the greater solubility in the hydrocarbon suggest about the bonding in iodine?
The colour of the gas disappears / Solution turns colourless to dark brown
(iii) The percentage composition of the elements in the solvent is:
element
% composition by mass
carbon
92.3
hydrogen
7.7
Calculate the empirical formula of the solvent.
(c) If the hydrocarbon layer is shaken with aqueous sodium thiosulphate the colour disappears.
(i) Write an equation for the reaction between iodine and sodium thiosulphate.
(ii) If insufficient thiosulphate solution were added traces of iodine are left which are not visible to the eye. Name a reagent that could be added to detect the iodine and state the colour produced.
Section A
For each of the following questions only one of the lettered responses (A-D) is correct.
Select the correct response in each case and mark its code letter by connecting the dots as illustrated on the response sheet.
1.B Which one of the following equations does not represent a reaction of chlorine under suitable conditions?
A Cl2 + H2O HCl + HOCl
B 2Cl2 + CH4 CCl4 + 2H2
C 3Cl2 + 6NaOH 5NaCl + NaClO3 + 3H2O
D Cl2 + 2NaOH NaCl + NaOCl + H2O
2.A When the hydrogen halides dissolve in water acidic solutions are formed. Which has the highest pH (assuming all are equal concentrations)?
A Hydrogen fluoride
B Hydrogen chloride
C Hydrogen iodide
D Hydrogen bromide
3. In which one of the following changes has chlorine been oxidised?
A 3Cl2 + 2Fe 2FeCl3
B Cl2 + I2 2IC1
C Cl2 + 2KBr Br2 + 2KCl
D Cl2 + 2KOH H2O + KClO + KCl
4. In which one of the following sequences are the oxides of the elements classified as basic, amphoteric and acidic respectively?
A Na, Mg, Al
B Na, K, S
C K, Al, P
D S, P, Mg
5. Which one of the following is true of the halogens as the Group is descended from chlorine to iodine?
A The atomic radius decreases.
B The colour of the element lightens.
C The melting point of the element increases.
D The oxidising power of the element increases.
5. Which one of the following equations represents a reaction of chlorine?
A Cl2 + NaOH NaOCl + HCl
B Cl2 + Fe FeCl2
C Cl2 + H2O HCl + HOCl
D Cl2 + CH4 CH2Cl2 + H2
SECTION B
1. The group in the Periodic Table headed by fluorine is called the halogen group.
(a) Using your knowledge of the group predict the colour and physical state of astatine.
Black / solid
(b) When chlorine gas is bubbled through a solution of potassium iodide a redox reaction occurs.
(i) What is observed in this reaction?
The colour of the gas disappears / Solution turns colourless to dark brown
(ii) What is meant by a redox reaction?
A reaction in which one or more electrons / is~are transferred from one substance to another
(iii) Explain fully the changes undergone (if any) by the potassium and iodide ions.
Potassium ion : no change (1)/ iodide ion : loses electron to form iodine (1) [2]
(c)
(i) The presence of iodide ions in an aqueous solution of sodium iodide can be detected using silver nitrate solution and ammonia solution. Describe fully how you would carry out this procedure experimentally, stating any changes you observe. Write equations for any reactions which occur.
Add dilute nitric acid followed by silver nitrate solution to the iodide solution/ A pale yellow / precipitate / of silver iodide / forms. This is insoluble in ammonia solution. / AgNO3 + NaI = AgI + NaNO3 (1)
(ii) State what differences you would observe if the sodium iodide solution was replaced by sodium chloride solution. Write equations for any reactions which occur.
With sodium chloride a white precipitate of silver chloride forms. This is soluble in ammonia solution
2. Sodium bromide reacts with concentrated sulphuric acid at room temperature to give four products. Place ticks (√) in four of the boxes to show which substances are formed.
√ if the substance is formed
bromine
hydrogen
hydrogen bromide
hydrogen sulphide
sodium hydrogensulphate
sulphur
sulphur dioxide
3. The diagram below shows chlorine gas being passed through a dilute
solution of potassium iodide. The upper layer is a hydrocarbon solvent.
(a) (i) What is the most important safety precaution to take when carrying
out the experiment?
The colour of the gas disappears / Solution turns colourless to dark brown
(ii) Write the ionic equation for the reaction between chlorine and potassium iodide and explain the redox reactions taking place in terms of electron transfer.
With sodium chloride a white precipitate of silver chloride forms. This is soluble in ammonia solution
(b) When the aqueous layer is shaken with the hydrocarbon most of the iodine dissolves in the upper layer.
(i) What is the colour of iodine in potassium iodide solution?
(ii) What does the greater solubility in the hydrocarbon suggest about the bonding in iodine?
The colour of the gas disappears / Solution turns colourless to dark brown
(iii) The percentage composition of the elements in the solvent is:
element
% composition by mass
carbon
92.3
hydrogen
7.7
Calculate the empirical formula of the solvent.
(c) If the hydrocarbon layer is shaken with aqueous sodium thiosulphate the colour disappears.
(i) Write an equation for the reaction between iodine and sodium thiosulphate.
(ii) If insufficient thiosulphate solution were added traces of iodine are left which are not visible to the eye. Name a reagent that could be added to detect the iodine and state the colour produced.
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