The first three periods

1
The first three periods
2
Periodic variation of the properties of elements
3
Structures of Period 3 Elements
4
Expanding the Octet

• Nitrogen cannot form five covalent bonds because
  it has no low-lying d-orbitals.
• The approach of an electronegative atom assists
  the excitation of the outer 3s electrons in
  P(3s23p3) to the vacant 3d orbitals, forming
  P(3s13p33d1).
• In Phosphorus, the promotion of an electron from
  3s to 3d requires 1370 kJ mol-1.
• Being excited, P(3s13p33d1) forms 5 covalent
  bonds with the electronegative atoms. Despite the
  use of energy in electronic promotion, the
  formation of PCl5 is favored because the
  formation of five P-Cl bonds releases energy.

5
Periodicity in behavior of chlorides with water

• The first three period 3 chlorides are ionic.
  MgCl2 slightly hydrolyzes to give a weakly acidic
  solution
• AlCl3 hydrolyzes to form an acidic solution. Al3
  in AlCl3 increases the covalent character in
  AlCl3
• SiCl4, PCl3 and S2Cl2 are covalent liquid
  chlorides which react with water to form HCl
  fumes.
• In going across a Period, as the bonding of the
  chlorides changes from ionic to covalent, the
  chlorides become increasingly acidic, hydrolyzing
  in water to give hydrogen ions.
• CCl4 is a covalent chloride which does not
  hydrolyze
• SiCl4 H2O Si(OH)4 4HCl

6
Periodicity in the behavior of hydrides with water

• Periods 2 3 hydrides are covalent except LiH
• Ionic hydrides react with water to give hydrogen
  and an alkaline solution.
• In going across a Period, as the bonding of the
  hydrides changes from ionic to covalent, the
  hydrides hydrolyze in water to produce an
  alkaline then neutral and eventually acidic
  solution toward the end of the Period.
• LiH H2O LiOH H2
• HF H2O H3O F-
• H2S H2O H3O HS-
• Group 5 hydrides such as NH3 PH3 give basic
  solution

7
Periodicity in the behavior of oxides with water

• Period 3 oxides are solids except SO2.
• In going across a Period, as the difference in
  electronegativity between the element oxygen
  decreases, bonding in oxides changes from ionic
  through partial ionic partial covalent to
  predominantly covalent. Corresponding to this
  change, the oxides also change from strongly
  basic, through amphoteric to strongly acidic.
• Li2O 2H2O 2LiOH H2
• BeO is amphoteric it doesnt react with
  water(Al2O3)
• B2O3 3H2O 2H3BO3 (very weak acid)
• 2NO2 H2O HNO2 HNO3
• OF2 H2O 2HF O2

8
Reactions of oxides

• SiO2 does not react with water, but it is acidic
• Heavier Period 3 elements form covalent oxides,
  which give acidic solution
  P4O10 6H2O 4H3PO4
  SO3 H2O
  H2SO4 Cl2O7 H2O
  2HClO4
• As the structure bonding of the oxides change
  from giant ionic to simple molecular across a
  Period, their acid-base nature also changes
  correspondingly from basic, through amphoteric to
  acidic.

9
Diagonal Relationship

• Elements diagonally related in the Periodic Table
  usually exhibit similarities in their physical
  and chemical properties.
• Li and Mg form many insoluble salt their
  carbonates nitrates decompose readily on
  heating.
• Be and Al can both react with acids and alkalis
  Be(s) 2HCl(aq) BeCl2(aq) H2(g)
  Be(s) 2NaOH(aq)
  2H2O(l) Na2Be(OH)4(aq) H2(g)
• On going down the Periodic Table diagonally, one
  encounters elements with very similar cationic
  size and electronegativity, resulting in similar
  physical chemical properties. (BeCl2 AlCl3
  have similar mps.)
• BeCl2 and AlCl3 are intermediate chlorides with a
  high degree of covalent character.

10
Effects of bonding on hydrolysis

• Group 1 chlorides are ionic due to the large
  difference in electronegativity of the elements.
• The strong non-directional electrostatic
  attractions between the ions result in their high
  melting points
• In MgCl2, the Mg2 polarizes the Cl- to some
  extent This decreases the degree of ionic
  character and changes somewhat its chemical
  behavior. The high covalent character of the
  Li-Cl bond accounts for the solubility of LiCl in
  ether and a hydrated form LiCl.2H2O.
• The attraction between Mg2 and the oxygen of H2O
  is great enough that the O-H bond in H2O is
  weakened. Eventually one of the surrounding water
  molecules has its O-H bond broken to result in
  the partial hydrolysis of MgCl2(s). Hydrated
  MgCl2(s) warm H2O MgCl(OH)(s) HCl(aq)

11
Intermediate chlorides

• In BeCl2, the electron cloud is so polarized that
  its bonding is essentially covalent.
• The high covalent character in BeCl2 results from
  an incomplete transfer of charge from Be to Cl
• When BeCl2 solid melts, polarization contributes
  to the stability of the ion pairs formed and the
  liquid state has an extra stability relative to
  the solid, with the result that melting point for
  BeCl2 is lower.
• BeCl2 exists as a linear molecule in the gaseous
  state. In the solid state it is polymeric,
  consisting of an infinite assembly of chains. In
  its solid state , AlCl3 consists of a layer
  lattice. The vapor contains Al2Cl6 molecules.
• The sp3 hybridization results in aluminum having
  four tetrahedrally arranged sp3 orbitals. In
  Al2Cl6 molecule, there are 3 normal covalent bond
  pairs and one dative covalent bond with Cl
  donating 2 electrons.

12
Bonding in Aluminum chloride

• In liquid state Al2Cl6 dimers exist. In gaseous
  state the dimer dissociates at high temperature
  into its monomers Al2Cl6(l) 2AlCl3(g)
• The AlCl3 monomer has a vacant p orbital that is
  perpendicular to the plane of the molecule. An
  Al2Cl6 dimer can be considered as being formed
  when a Cl atom attached to one Al atom donates a
  pair of electrons to the neighboring Al atom,
  thus giving each Al atom a noble gas
  configuration and a tetrahedral arrangement of
  Cl atoms around it.
• Due to its high covalent character, AlCl3
  dissolves in ethanol, forming a Lewis acid
  catalyst.

13
Hydrolysis of covalent chlorides

• BeCl2 AlCl3 still retain some ionic character
  and dissolve readily in aqueous solution. Owing
  to the small size high charge of the Be2
  Al3 ions, they are solvated and undergo vigorous
  hydrolysis, liberating H ions (Acidic solution).
• Be(H2O)42(aq)H2O(l) Be(OH)(H2O)3(aq)H3O
  (g)
• Solid Aluminum chloride fumes in moist
  air(hygroscopic) 2AlCl3(s) 3H2O(l)
  Al(OH)3(s) 3HCl(g)
• When dissolved in water, aluminum chloride gives
  2AlCl3(s) 6H2O(l) Al(H2O)63(aq)
  3Cl-(aq) Al(H2O)63(aq)H2O(l)
  Al(H2O)5(OH)2(aq)H3O(aq)
• 6 water molecules form dative bonds with the Al3
  because the Al atom expands its octet by using
  the 3d subshell.

14
Properties of AlCl3(aq)

• Adding OH- or CO32- to aqueous solution of AlCl3
  causes loss of H3O, shifting the equilibrium
  position. Proton loss continues until hydrated
  Al(H2O)3(OH)3 (aluminum hydroxide) is formed.
• Al(OH)(H2O)52(aq) OH-(aq)
  Al(OH)2(H2O)4(aq) H2O(l) Al(OH)2(H2O)4(aq)
  OH-(aq) Al(OH)3(H2O)3(s) H2O(l)
• Adding excess OH- forms hexahydroxoaluminate(III)
  ion Al(OH)63-.
• Al(OH)3(H2O)3(s) is amphoteric as it takes up
  OH-/H3O ions.
• The m.p. (190oC) and b.p.(420oC) show that
  aluminum chloride is covalent in character.
  Solubility in benzene shows that the compound is
  non-polar in nature. The relative molecular mass
  of 267 indicates that the compound exists as
  dimer in benzene.
• AlCl3 hydrolyzes in aqueous solution to give an
  acidic solution and the acidity is due to the
  Al3 ion.

15
Properties of Period 2 3 Covalent Chlorides

• The central atom undergoes different types of
  hybridization for bonding with the Cl-atom.
• These covalently bonded molecules are usually
  liquids or gases with low m.p and low b.p.
• They do not conduct electricity dissolves in
  organic solvents.
• BCl3 hydrolyzes vigorously in water
  BCl3(l) 3H2O(l) H3BO3(aq) 3HCl(g)
• The B-Cl bond is polarized with the sp2
  hybridized boron atom highly electropositive. The
  vacant p orbital perpendicular to the plane of
  the BCl3 molecule thus accepts readily the lone
  pair on the O-atom in the attacking water
  molecule. A HCl molecule is quickly eliminated
  from the tetrahedral intermediate to restore the
  planar molecular structure. This process is
  repeated until all the -Cl groups are replaced by
  -OH groups.

16
Group IV Covalent Chlorides

• The non-polar CCl4 molecule is inert as there are
  no low-lying vacant orbitals for attack of water
  molecule.
• SiCl4 hydrolyzes readily in moist air, producing
  misty fumes of hydrogen chloride an a white
  precipitate of hydrated oxide
  
  SiCl4(l)4H2O(l) Si(OH)4(l)4HCl(g)
  SiO2.2H2O(s)
• The Si-atom in SiCl4 has low energy d orbitals
  available for holding the lone pair of oxygen in
  the attacking water molecule, giving rise to a
  five-coordinated intermediate. A HCl molecule
  is easily eliminated from the intermediate.
• The Si-atom in SiCl4 is sp3 hybridized (4 sp3
  hybrid orbitals). There are vacant 3d orbitals in
  the Si-atom which are slightly higher in energy.
  These vacant orbitals accept the lone pair from
  water during the hydrolytic attack.

Silica seals the stopper
17
Bonding and Structure of Covalent Chlorides

• 4 factors affecting the hydrolysis of covalent
  chlorides
• Steric hindrance
• Bond polarity
• Enthalpy reasons
• Availability of low energy d-orbitals
• SiCl4 dissolves in conc. HCl
• SF6 is as unreactive as nitrogen.
• NCl3 attacks H2O with its lone pair, forming HOCl
• The central atom in both SiCl4 PCl3 are
  electron deficient and this affects its mechanism
  of attack by the water molecules. Like SiCl4, the
  vacant orbitals in phosphorus can hold the lone
  pair of oxygen in the attacking water molecules,
  resulting in hydrolysis in a different manner.

18
Mechanism of hydrolysis of some covalent chlorides

• The energy given out on the formation of a PO
  bond is the driving force for the hydrolysis of
  bot PCl3 PCl5.

PCl3(l) 3H2O(l) H3PO3(aq) 3HCl(aq)
19
Group VI Covalent Chlorides

• Like NCl3, the O-atom in OCl2 has no low energy d
  orbital available and it undergoes hydrolysis by
  using its lone pair of electrons

The hydrolysis of SCl2(g) and S2Cl2(g) is similar
SCl2(g) H2O(l) HSCl(aq)
HOCl(aq) 2S2Cl2(g) 2H2O(l) 3S(s)
SO2(g) 4HCl(aq)
20
Period 2 and Period 3 Hydrides
Ionic Hydrides
Covalent hydrides with some ionic character
Polar Covalent Hydrides
Typical Covalent Hydrides
There is a change in hydrolytic properties across
a Period with the change in nature of bonding of
the hydrides.
The hydride ion H- is an even more powerful base
than the OH-. When dissolved in aqueous solution
it abstracts a H readily from water. H-
H2O H2 OH-
21
Reactions of Covalent Hydrides

• Group II hydrides hydrolyse readily in aqueous
  solution MH2(s) 2H2O(l) M(OH)2(aq)
  2H2(g)
• They form a less alkaline solution than the Group
  I hydrides, as their hydroxides M(OH)2 are less
  soluble because of their higher lattice
  enthalpies.
• Of the Group III hydrides, B2H6 is an inflammable
  gas which is decomposed readily by water
  B2H6(g) 6H2O(l)
  2H3BO3(aq) 6H2(g)

AlH3(g) 3H2O(l) Al(OH)3(aq)
3H2(g)
Owing to its small size, the C-atom can form
stable covalent bonds with other carbon atoms.
The resulting C-C bond has such a high intrinsic
strength (bond enthalpy 356 kJ mol-1) that it
is almost as strong as the CO bond (bond
enthalpy 360 kJ mol-1). This makes oxidative
decomposition of carbon compounds to such
products as CO2 H2O less energetically favored.
Hydrides of carbon are thus generally fairly
stable.
22
Groups IV and V Hydrides

• Si-O bonds (bond enthalpy 464 kJ mol-1) are
  much stronger than Si-Si bonds(bond enthalpy230
  kJ/mol) because of the smaller size of the
  O-atom. Thus, Si-Si bonds are easily oxidized to
  Si-O bonds (whereas C-C bonds are not), and
  catenated hydrides of silicon are energetically
  less stable with respect to (SiO2)n
• SiH4(g) 2O2(g) SiO2(s) 2H2O(g)
• Because of its small size the C-atom forms ?
  bonds by sideways overlapping of p orbitals thus
  formation of hydrocarbons with CC, CC and C..C
  bonds are possible. Because of the larger size,
  Si cannot form double and triple bonds.
• The P-H bond in PH3 involves the overlap of an
  sp3 orbital from the 3rd shell of Phosphorus,
  with that of an s orbital from the 1st shell of
  hydrogen. The sigma overlap between orbitals of
  widely different energies will be ineffective,
  resulting in weak P-H bonds.
• Formation of PH5 is not favored as the 5 P-H
  bonds are weak
• PH3 is more reactive to oxidation owing to the
  formation of the much more stable product P2O5
  with high P-O bond energies

23
Polar Covalent Hydrides

• The order of reducing strength PH3gtH2SgtNH3gtH2O
  Across Periods 2 3, there is a decrease in
  atomic size and an increase in electronegativity.
  The loss of electrons from S-atom and O-atom is
  more difficult, so that H2O and H2S are weaker
  reductants.
• Period 3 hydrides such as PH3 H2S are more
  reactive toward oxidation because P and S can
  expand the octet, exhibiting more oxidation
  states.
• Higher oxidation states are favored by more
  electropositive atoms and the drive for
  phosphine to achieve the stability of 5 state in
  P2O5 is particularly strong.
  2PH3(g) 4O2(g) P2O5(s)
  3H2O(l)
• Besides PH3 has a positive /\Hf of 5.4 kJ mol-1
  and is expected to be very unstable. /\Hf
  H2O(-286 kJ/mol), NH3(-46 kJ/mol)

24
Group I II oxides/hydroxides

• Group I II oxides have ionic lattices and are
  highly basic they react exothermically with
  water to give the hydroxides. For example,
  Na2O2(s) H2O(l) 2NaOH(aq)
  ? O2(g) Li2O(s) H2O(l)
  2LiOH(aq) CaO(s) H2O(l)
  Ca(OH)2(aq)
• Group I hydroxides are all soluble and the base
  strength increases down the group. The solubility
  base strength of Group II hydroxides increase
  down the group.
• Be(OH)2 is insoluble amphoteric Ba(OH)2 is
  strong alkali.
• The s-block elements form compounds that are
  predominantly ionic in nature, showing constant
  oxidation numbers of 1, 2.
• Energy released in Cl(g)e---gtCl-(g)
  Na(g)Cl-(g)--gtNaCl(s) more than covered the
  energy used in atomization (Na(s)-gtNa(g), ?
  Cl2(g) --gt Cl(g)) 1st ionization of Na(g).

25
Weak tendency to form complexes

• A complex is a polyatomic ion/neutral molecule
  formed when molecular/ionic groups (ligands) form
  dative covalent bonds with a central metal
  atom/cation.
• The low energy d-orbitals in the d-block metal
  ions accept the lone pairs from the surrounding
  ligands (Cl-, NH3, OH2), forming dative covalent
  bonds.
• S-block metal ions may be surrounded by polar
  molecules, with the negative ends of their
  dipoles towards the cation. This is called
  solvation or hydration if the solvent is water.
  The association is just the electrostatic
  attraction between the dipoles and the cation.
  S-block metal ions have no low energy vacant
  orbitals available for bonding with the lone
  pairs of surrounding ligands they do not form
  complexes.

26
Variation in properties of the s-block elements

• Lithium has the smallest size among the Group I
  elements. When it forms the Li ion, the
  remaining electrons experience the strongest
  nuclear attraction as they do not screen each
  other from the full nuclear attraction.
• There is a general increase in ionic atomic
  radius on descending both Groups I and II. Across
  each Period from Group I to II, the atomic and
  ionic radii decrease.
• Li Mg2 or Na Ca2 have similar ionic radii.

27
Variation in Ionization/Hydration Enthalpies
28
Hydration Enthalpies

• The hydration enthalpy is the energy released
  when H2O molecules cluster around 1 mole of metal
  ions M(g) aq M(aq)
• On going down both Groups I II, the charge
  density (charge-to-size ratio) decreases the
  coulombic energy of stabilization falls. As the
  ions get larger from Li to Cs, or from Be2 to
  Ba2, the electrostatic attraction between H2O
  molecules and the ions get less.

29
Variation in Melting Points

• The strength of the metallic bonds depends on
• Ionic radius,
• Number of electrons in the outer shell of an atom
• Metallic crystal structure (B.C.C.,H.C.P.,C.C.P)
• A metal can be considered as a fairly close
  packed lattice of cations cemented together by
  a delocalized sea of mobile valence electrons. As
  Groups I II are descended, the attractive
  forces between the electron sea and the ions
  become weaker because of the decreasing electron
  density the higher tendency of the electron sea
  to spread out. The general decrease in melting
  points on descending Groups I II is a
  reflection of the decrease in strength of their
  metallic bonds.
• The greater number of valence electrons and the
  smaller ionic radii of the Group II elements help
  to increase the density of the electron sea
  thus increase the attractive forces between the
  ions and the electron cloud.

30
Reactions of Groups I II metals in aqueous
solution

• Eo is a measure of the process M(aq) e-
  M(s)
• The highly negative values for Groups I II
  imply that the reverse process is more likely to
  occur.
• These metals are good reductants, having a strong
  tendency to form cations in solutions.
• Li(aq) has the most negative Eo because of the
  unusually small size of Li ion. This makes the
  cation heavily hydrated, resulting in a highly
  negative enthalpy of hydration.
• For Group 2 elements, the positive contribution
  of Eo, due to the sum of the 1st and 2nd
  ionization enthalpies, are even greater than the
  negative contribution due to the hydration
  enthalpy, when they are compared with the Group I
  elements.
• Berylliums exceptionally less negative Eo value
  is the result of its high ionization enthalpies
  atomization enthalpy. Beryllium does not react
  with water at all.

31
Reactions of Groups I II metals

• State functions like Eo and /\H all indicate that
  reaction of lithium with water is feasible. The
  higher m.p. of Li increases the activation energy
  required for dissolution in aqueous solutions,
  decreasing its reducing power.
• Group I metals burn brilliantly in air to form
  one or more of the three types of oxide. Sodium
  produces a mixture of sodium oxide sodium
  peroxide. Potassium, rubidium caesium react
  with oxygen to produce superoxides, containing
  the O2-. Like Li, M2 polarizes the larger O2-
  ions do not form superoxides.
• All the s-block metals except Be react directly
  with H2. 2Na(s) H2(g) /\ 2NaH(s)
  The reactivity
  increases with increasing reducing power down
  Groups I II. S-block hydrides (except BeH2 and
  MgH2) are ionic H- could be discharged as H2 at
  anode.

32
Variation in properties of the s-block compounds

• All Group I oxides react exothermically with
  water to form hydroxides Li2O(s) H2O(l)
  2LiOH(aq)
• Peroxides give the metal hydroxide hydrogen
  peroxide Na2O2(s) 2H2O(l)
  2NaOH(aq) H2O2(aq)
• Superoxides give the metal hydroxide, hydrogen
  peroxide oxygen 2KO2(s)2H2O(l)
  2KOH(aq) H2O2(aq)O2(g)
• Both peroxides and superoxides are oxidizing
  agents.
• KO2 is used for purifying the air in submarines,
  gas masks. 4KO2(s) 2CO2(g)
  2K2CO3(s) 3O2(g)
• Except LiOH, Group I hydroxides are strongly
  basic absorb CO2 to form carbonates or hydrogen
  carbonates 2OH-(aq) CO2(g)
  CO32-(aq) H2O(l)
  OH-(aq) CO2(g) HCO3-(aq)

33
Group II oxides/peroxides

• Group II oxides are less basic than Group I
  oxides. Like Al2O3, BeO has a high degree of
  covalent character and is amphoteric BeO(s)
  2H(aq) Be2(aq) H2O(l) BeO(s)
  2OH-(aq) H2O(l) Be(OH)42-(aq)
• Basicity of Group II oxides increases down the
  Group.
• The solubilities of Group II hydroxides increase
  down the Group. Due to the high solubility of
  Ba(OH)2, H2O2 cannot be prepared from BaO2. To
  prepare H2O2 in the laboratory, barium peroxide
  is added to ice-cold dilute H2SO4. BaSO4 is
  separated from H2O2 by filtration. BaO2(s)
  2H2O(l) Ba(OH)2(aq) H2O2(aq)
  BaO2(s) H2SO4(aq) BaSO4(s)
  H2O2(aq)

34
S-block chlorides/hydrides/hydroxides/carbonates

• Group I chlorides are anhydrous except LiCl.
  Unlike Group II chlorides easily react with water
  to give hydrates. Anhydrous CaCl2 is
  deliquescent. The formula for the hydrates
  MgCl2.6H2O, BaCl2.2H2O
• s-block hydrides readily react with water to give
  the metal hydroxide and hydrogen, the vigor of
  reaction increasing down the Group.
  H-(aq)H2O(l) OH-(aq)H2(g)
• Hydrides in the form of LiAlH4 NaBH4 are used
  in organic chemistry as selective reductants
  which reduce CO, not CC
• Group II carbonates and hydroxides are thermally
  unstable and thermal stability increases on
  descending Groups I II.
• For compounds with large polarizable anions
  (CO32-, OH-), the thermal stability increases as
  the polarizing power of the cation decreases. Due
  to the strong polarizing power of Li, Li2CO3 has
  a higher covalent character and is liable to
  decomposition to the smaller oxide anions, so
  that the attraction between the ions become
  greater.

35
Stability of Group I II oxides

• 2LiOH(s) Li2O(s) H2O(g)
  Li2CO3(s) Li2O(s)
  CO2(g)
• The O2- is less readily polarized by the cation
  and its higher charge density makes it more
  strongly attracted to the cation, thus enabling
  greater lattice stability. The lattice enthalpies
  of the oxides decrease in the order Li2O gt Na2O2
  gtKO2 gtRbO2 BeO
  gt MgO gt CaO gt SrO gt BaO
• The change in lattice enthalpy from Li2CO3 to
  Li2O, is thus larger than te change from Na2CO3
  to Na2O2, thus making the decomposition of Li2CO3
  energetically more favorable and hence more
  liable to take place.
• The large charge density in the M2 makes the
  M2O lattice more stable than the M2O or M2O2
  lattice, thus energetically favoring
  decomposition of M2CO3 more.

36
Relative solubility of sulphates and hydroxides

• The solubility of the hydroxides increases down
  Group II the solubility of the sulphates
  decrease down Group II.
• If an ionic solid is to dissolve in a solvent
  there must be energetically favorable
  interactions between the solvent and the
  dissolved ions to compensate for losing the
  favorable solute-solute interactions in the ionic
  solid.
• Strong ion-dipole interactions account for high
  solubility

37
Solubilities of hydroxides sulphates

• Ions in an ionic solid are separated to such a
  distance that there is no longer any interaction
  between them, as is nearly the case in solutions
  with infinite dilution. NaCl-(s) aq
  Na(g) Cl-(g) This
  process is just the reverse of lattice formation
  and requires an enthalpy which is equal to the
  lattice energy with the sign reversed.
  /\Hinfinite dilution 776 kJ mol-1
• NaCl(s) Na(g) Cl-(g) -/\HlatticeNaCl)
  776 kJ mol-1
• The enthalpy change that accompanies the
  hydration of one mole of both of these gaseous
  ions is called the enthalpy change of hydration
  Na(g)Cl-(g)
  aq Na(aq)Cl-(aq) /\Hhydration-772
  kJ/mol
• The enthalpy change when 1 mole of NaCl(s) is
  dissolved in such an amount of water that further
  dilution gives no detectable heat change
  (infinite dilution) is 4 kJ/mol

38
Relative solubility of hydroxides and sulphates

• Due to its lower charge density, /\Hlattice of
  Group I compounds are smaller than those of Group
  II compounds. Thus /\Hsolution for Group I
  compounds is more exothermic.
• For the sulphates of Group II metals, the cations
  are much smaller than the anions. Down the Group,
  the change in size of the cations does not cause
  a significant change in /\Hlattice.
• The changing size of the cations cause
  /\Hhydration to become less exothermic down the
  Group. Thus, /\Hsolution becomes less exothermic
  and the solubility of Group II sulphates
  decreases down the Group.
• For the hydroxides of Group II metals, the sizes
  of the anions cations are of the same magnitude
  and less enthalpy is required to break the
  lattice as the cationic size increases down the
  group. (/\Hlattice ? 1/(r r-))
• As the change in /\Hhydration is small, the
  /\Hsolution becomes more exothermic the
  solubility of Group II hydroxides increases down
  the Group.

39
Reaction of Groups I II with Chlorine

• The enthalpy change of formation of Group I
  metallic chlorides rises as the Group is
  descended and the ionization enthalpies of the
  metallic elements fall.
• Due to its high ionization enthalpy, Be forms
  covalent chloride, with comparatively low m.p.

40
Group VII hydrides

• A 0.1 M HF(aq) is a much weaker acid than that of
  a 0.1M HCl(aq). The very large H-F bond energy
  makes the enthalpy of hydration of F-
  insignificant.
• The dissociation of H-F occurs to a lesser
  extent. HF(aq) H2O(l) F-(aq)
  H3O(aq) HCl(aq) H2O(l)
  Cl-(aq) H3O(aq)