1. Lecture 12. S-elements of the ІІ А group. Alkaline earth metals
Lecture 12. S-elements of the ІІ А group. Alkaline earth metals. р-Elements of the ІІІА group. Boron and Aluminium Be Ra Mg Ba Sr Ca
PhD Halina Falfushynska

2. Members of the s-Block ElementsLi Be Na K Rb Cs Fr Mg Ca Sr Ra Ba
IA IIA
IA Alkali metals
IIA Alkaline Earth
metals

3. High tendency to lose e- to form positive ions
Metallic character increases down both groups
Group II Be
Mg 1.2 Ca Sr Ba Ra
Low nuclear attraction for outer electrons
Highly electropositive
Small electronegativity

4. Occurrence and Extraction
These elements are widely distributed in rock structures. The main minerals in which magnesium is found are carnellite, magnesite and dolomite.
Calcium is found in chalk, limestone, gypsum and anhydrite.
Magnesium is the eighth most abundant element in the Earth’s crust, and calcium is the fifth.
It is extracted from sea-water by the addition of calcium hydroxide, which precipitates out the less soluble magnesium hydroxide. This hydroxide is then converted to the chloride, which is electrolysed in a Downs cell to extract magnesium metal. MgCl2 → Ca + Cl2;
cathode: Mg2+ + 2e → Mg0 ; anode: 2Cl- – 2e → Cl02
Beryllium: BeF2 + Mg  MgF2 + Be.

5. Biological occurrence of elements of group IIA
Beryllium's low aqueous solubility means it is rarely available to biological systems, is usually highly toxic.
Magnesium and calcium are ubiquitous and essential to all known living organisms. They are involved in Mg/Ca ion pumps, magnesium functioning as the active center in some enzymes, and calcium salts taking a structural role (e.g. bones).
Strontium and barium have a lower availability in the biosphere. Strontium plays an important role in marine aquatic life, especially hard corals. They use strontium to build their exoskeleton. These elements have some uses in medicine, for example "barium meals" in radiographic imaging, whilst strontium compounds are employed in some toothpastes.

6. Flame test Mg brilliant white Ca brick red Sr blood red Ba apple green
HCl(aq)
sample

7. Basic oxides, hydroxides
BeO
Be(OH)2
MgO
Mg(OH)2
CaO
Ca(OH)2
SrO
Sr(OH)2
BaO, Ba2O2
Ba(OH)2
Amphoteric to basic, base strength increase
Reaction with water:
Oxide: O2- + H2O  2OH-

8. Predominantly ionic with fixed oxidation state
Group II: Electropositive metals.
Low first and second I.E. but very high third
I.E Have a fixed oxidation state of +2.
Be and Mg compounds possess some degree
of covalent character.
Compare to
Group I: Most electropositive metals.
Low first I.E. and extremely high second I.E.

9. Atomic radii (nm) Li 0.152 Be 0.112 Na 0.186 Mg 0.160 K 0.231 Ca 0.197Rb
0.244 Sr
0.215 Cs
0.262 Ba
0.217 Fr
0.270 Ra
0.220 Fr Ra Li Be

10. Ionization Enthalpy Li Na K Rb Cs 1st I.E. 300 400 500 600 500 1000
1500
2000 Be Ca Ba
Be+
Ca+
Ba+
1st IE
2nd IE

11. Variation in Melting Points
250
500
750
1000
1250 Be Mg Ca Sr Ba Li Na K Rb Cs

12. Reactions with oxygen Normal Oxide Peroxide Superoxide Formed by
S-block elements are strong reducing agents.
Their reducing power increases down both groups.
(As the atomic size increases, it becomes easier to
remove the outermost electron)
S-block elements reacts readily with oxygen.
they have to be stored under liquid paraffin
to prevent contact with the atmosphere.
Normal Oxide
Peroxide
Superoxide
Formed by
Be, Mg, Ca, Sr Ba
None

13. Reaction with hydrogen
Reaction with water
Reaction with hydrogen
M(s)  M+(aq) + e-
H2O(l) + e-  OH-(aq) + ½ H2(g)
All the II A elements except Be react directly with
hydrogen.
Ca(s) + H2(g)  CaH2(s)
The reactivity increases down the group.
Only BeH2 and MgH2 are covalent, others are ionic.
Li volt Na K Rb Cs
Be volt
Mg -2.38 Ca Sr Ba

14. Reaction with chlorine
Reactions of chlorides
Group II chlorides show some degree of covalent character.
BeCl2 hydrolysis to
form Be(OH)2(s) and HCl(aq).
MgCl2 is intermediate, it dissolves and
hydrolysis slightly.
Other group II chlorides just dissolve without
hydrolysis.
All the s-block metals react directly with chlorine to produce chloride.
Ba (s) + Cl2(g)  BaCl2 (s)
BeCl2 is essentially covalent
The lower members in group II form essentially ionic
chlorides, with Mg having intermediate properties.

15. Reaction with nitrogen
Reactions with acid
All the IIA metals react directly with acid to produce salt and hydrogen.
Mg (s) + H2SO4(l)→ MgSO4 (s) + H2 (g)
Mg + 2HCl -----> MgCl2 + H2
These reactions can not occur without extreme circumstances.
A compound may be created via really high temperatures.
3Mg(s) + N2(g) -> Mg3N2(s)
Reactions with alkali. Only for Beryllium!!!
Be (s) + 2NaOH(l) + 2H2O = Na2[Be(OH)4] (s) + H2

16. Reactions of oxides and hydroxides
All group I oxides reacts with water to form
hydroxides
Oxide: O2- + H2O  2OH-
Peroxide: O H2O  H2O2 + 2OH-
Superoxide: 2O2- + 2H2O  2OH- + H2O2 + O2
All group I oxides/hydroxides are basic and the
basicity increases down the group.

17. Thermal Stability of carbonates
BeCO3  BeO + CO2 ( at 100oC)
MgCO3  MgO + CO2 ( at 540oC)
CaCO3  CaO + CO2 ( at 900oC)
SrCO3  SrO + CO ( at 1290oC)
BaCO3  BaO + CO2 ( at 1360oC)
Li2CO3  Li2O + CO2 ( at 700oC)
All other group I carbonates are stable at ~800oC

18. Thermal Stability of hydroxides
Be(OH)2(s)  BeO(s) + H2O(g) H = +54 kJ/mol
Mg(OH)2(s)  MgO(s) + H2O(g) H = +81 kJ/mol
Ca(OH)2(s)  CaO(s) + H2O(g) H = +109 kJ/mol
Sr(OH)2(s)  SrO(s) + H2O(g) H = +127 kJ/mol
Ba(OH)2(s)  BaO(s) + H2O(g) H = +146 kJ/mol
All group I hydroxides are stable except LiOH
at Bunsen temperature.

19. Explanation of Thermal Stability- +
Decreasing
polarizing
power
Increasing
stability + - + -

20. Relative solubility of Group II hydroxides
Compound
Solubility / mol per 100g water
Mg(OH)2
0.020 x 10-3
Ca(OH)2
1.5 x 10-3
Sr(OH)2
3.4 x 10-3
Ba(OH)2
15 x 10-3
Solubility of hydroxides
increases down the group.

21. Solubility of Group II sulphates
Compound
Solubility / mol per 100g water
MgSO4
epsom salt
3600 x 10-4
CaSO4
11 x 10-4
SrSO4
0.62 x 10-4
BaSO4
0.009 x 10-4
Solubility of sulphates
increases up the group.

22. Explanation of solubility
Group I compounds are more soluble than Group II
because the metal ions have smaller charges and
larger sizes. H lattice is smaller, and H solution is
more exothermic.
H solution
-H lattice
H hydration = +

23. For Group II sulphates, the cations are much smaller
than the anions. The changing in size of cations does
not cause a significant change in H lattice (proportional
to 1/(r+ + r-).
However, the changing in size of cations does cause
H hydration (proportional to 1/r+ and 1/r-) to become less
exothermic, and the solubility decreases when
descending the Group.
SO42-
MgSO4
SrSO4

24. For the smaller size anions, OH-.
Down the Group, less enthalpy is required to break the lattice as the size of cation increases. However the change in H solution is comparatively smaller due to the large value of 1/r- .
As a result, H solution becomes more exothermic
and the solubility increases down the Group.
Mg(OH)2
Sr(OH)2

25. Comparison of the physical properties and chemical properties between alkaline earth and alkali metals
---The main difference is the electron configuration, which is ns2 for alkaline earth metals and ns1 for alkali metals. For the alkaline earth metals, there are two electrons that are available to form a metallic bond, and the nucleus contains an additional positive charge. Also, the elements of group 2A (alkaline earth) have much higher melting points and boiling points compared to those of group 1A
---The alkaline earth metals are much harder and denser compare to alkali metals.
---The elements of the group 2A contain a smaller atomic radius and much higher ionization energy than the group 1A. Even though the group 2A contains much higher ionization energy, they still form an ionic compound with 2+ cations.

26. Uses of IIA group compounds
Magnesium
Magnesium hydroxide - Milk of magnesia, an antacid
Magnesium can be obtained by eating food which is rich in Magnesium, such as nuts and certain vegetables or by eating supplementary diet pills. 
Calcium hydroxide
To neutralize acids in waste water treatment
Strontium compound
Strontium ranelate is used in the treatment of osteoporosis. It is a prescription drug in the EU, but not in the USA.
Strontium chloride is sometimes used in toothpastes for sensitive teeth. One popular brand includes 10% total strontium chloride hexahydrate by weight.
Barium
Uses of barium sulfate include being a radiocontrast agent for X-ray imaging of the digestive system

28. Properties and Trends in Group IIIA
Boron is a non-metal, whereas the other Group IIIA (13) elements are metals.
• The most common oxidation state for Group 13 is +3.
Gallium, indium, and thallium also often form 1+ ions by retaining their ns2 electrons; this is called the inert pair effect.
• Aluminium and gallium are amphoteric, but indium and thallium show more metallic character and do not dissolve in alkalis.

30. Production of Aluminum
The Hall–Heroult process is used to isolate aluminum.
Bauxite (aluminum ore) is first treated with NaOH to form aluminate ion, Al(OH)4–.
This solution is then diluted with water and acidified slightly, precipitating Al(OH)3. The Al(OH)3 is then heated to about 1200 oC and decomposes to pure Al2O3.
The Al2O3 is then electrolyzed using cryolite, Na3AlF6, as the electrolyte and graphite (solid carbon) as the electrodes.

32. Electrolysis Cell for Aluminum Production
Carbon
aluminium

33. Boron
Most of the chemistry of boron compounds is based on the lack of an octet of electrons about the central boron atom.
These compounds are electron-deficient; the deficiency causes them to exhibit some unusual bonding features.
Boron hydride (BH3) forms a coordinate covalent bond with another atom that has a lone pair of electrons to complete its octet; this is called an adduct.
Borax, Na2B4O7·10 H2O, a hydrated borate, is the primary source of boron.
Sodium perborate is used as a bleach; perborate’s structure is:

34. Structure of Diborane, B2H6
The “regular” B—H bond distance is shorter than the B—H distance in the three-center bond; why?

36. Properties and Uses of Aluminum
The reduction of Al3+(aq) to Al(s) occurs with difficulty.
Thus, aluminum metal is a good reducing agent.
As an active metal, aluminum readily reacts with acids to produce hydrogen gas.
Aluminum also dissolves in basic solutions.
Because its combustion is a highly exothermic reaction, powdered aluminum is used as a component in rocket propellants and fireworks.
Perhaps the most familiar use of aluminum is in beverage cans, cookware, and as a foil for wrapping foods.

37. Reactivity towards air
Boron is unreactive in crystalline form. Aluminium forms a very thin oxide layer on the surface which protects the metal from further attack. Amorphous boron and aluminium metal on heating in air form B2O3 and Al2O3 respectively. With dinitrogen at high temperature they form nitrides.

38. Oxides, hydroxides, acids
Boron forms the oxide B2O3 and a large number of borate anions containing trigonal planar BO3 units and/or tetrahedral BO4 units.
• Boric acid (B(OH)3) is a monobasic acid. It acts as a Lewis acid, interacting with water to form B(OH)3(OH2), which loses a proton to form [B(OH)4]–
• A12O3 and Ga2O3 are amphoteric, but In2O3 and T12O3 are basic.
• [M(H2O)6]3+ salts are acidic in aqueous solution due to hydrolysis.

39. Anodized Aluminum
The oxide layer on aluminum protects it from further corrosion.
The oxide layer may be enhanced by making an aluminum object the anode in an electrolysis apparatus.
The thick, hard oxide layer that results is porous and may be dyed.
When the pores are sealed, the color is a permanent part of the object.

40. Reactivity towards acids and alkalies
Boron does not react with acids and alkalies even at moderate temperature; but aluminium dissolves in mineral acids and aqueous alkalies and thus shows amphoteric character. Aluminium dissolves in dilute HCl and liberates dihydrogen. 2Al(s) + 6HCl (aq) → 2Al3+ (aq) + 6Cl-(aq) + 3H2 (g) However, concentrated nitric acid renders aluminium passive by forming a protective oxide layer on the surface. Aluminium also reacts with aqueous alkali and liberates dihydrogen

41. Reaction with halogens
Halides properties
The Lewis acidity of the boron halides increases in the order BF3 < BCl3 <
BBr3 < BI3. This is because it takes less energy to distort the trigonal planar
geometry with larger halides. This trend can be explained by weaker pπ–pπ bonding.
• Aluminium fluoride is a solid with high ionic character, but the other
aluminium halides have structures containing covalent bonds.
The IIIA elements react with halogens to form trihalides (except TlI3).
2E(s) + 3 X2 (g) → 2EX3 (s) (X = F, Cl, Br, I)