Group 2 elements and their compounds
Explain the reason why it happens towards the trend in physical properties of group 2 when down to group, discussion in relation to; atomic radius, first ionisation energy, electronegativity, melting/ boiling point, and atomisation energies. Give the reason why for each trend involved.
What happens when down group 2?
Progressing down group 2, the atomic radius increases due to the extra shell of electrons for each element.
1. Going down the group, the first ionisation energy decreases.
It is because increase in the size of the atom and the effective nuclear charge (+2) remains the same. The attraction between the nucleus and the valence electron gets weaker and is become easier to be removed. But the third ionisation energy for group 2 elements is exceptionally high as compared to the first and second ionisation energy. This is because the third electron is from a completely filled inner shell (8, octet) that is closer to the nucleus and experiences a greater attractive force as compared to the first two electrons. More energy than expected to required to remove the third electron.
Why M2+ not M+ or M3+ oxidation state for group 2 ?
It is because more negative Lattice Energy (???) more exothermic, due to the high charge density, high Lattice Energy, LE, high energy required. More stable the compound than M+ or M3+. All group 2 shows fixed oxidation state of +2 in all their compounds.
2. Going down the group, electronegativity decreases.
From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius.
3. Going down the group, melting point decreases.
The general melting point decreases as we move down the group, this is because it depends on the strength of the forces holding the particles together in the solid state as they are arranged in the solid lattice (lattice structure). All the group 2 elements are metals with strong metallic bonds holding the atoms together in the solid state, but same number of the electron delocalise, (2 e each atom for bonding) that donates for bonding.
When down to group 2, increase the proton number, increase size. The electrons are further from the nuclei of the increasingly large positive ions. This makes the metallic bonding weaker, reducing the melting temperature.
Lattice structure of group 2.
Be 1553K hexagonal close-packed, HCP
Mg 923K hexagonal close-packed, HCP
Ca 1113K face-centered cube, FCC
Sr 1041K face-centered cube, FCC
Ba 987K body-centered cube, BCC
Why the melting point of Mg is lower than Ca?
When down group 2, atomic radius increases, Mg is smaller size than Ca with the arrangement of HCP (more open) and FCC respectively. As a result, the metallic bond in magnesium is weaker than calcium. So that the melting point for magnesium is anomalously low. Refer to this discussion.
1. Going down the group, reactivity of Group 2 metals with water increases.
All alkaline earth metals react with water under suitable conditions to form their hydroxides or oxides and hydrogen. The overall trend, for the reactivity of Group 2 metals with water, is an increase down the group. The reactivity of the elements with water increases down the group due to the increase in the reducing strength of the metals (become easier to form cation, kekuatan reducing agent meningkat, reduced kan orang lain tetapi diri sendiri oxidised). Down group 2, the outer electrons are easier to remove as they are further from the nucleus and there is more shielding resulting in a lower nuclear attraction (1. far from nucleus, 2. low nuclear attract because of shielding effect).
More negative value of SREP, stronger the strength as a reducing agent.
2. Going down the group, solubility of _______ decreases / increases.
Solubility (ΔH solution, heat of solution) directly proportional to Heat of Hydration.
Solubility (ΔH solution, heat of solution) inversely proportional to Lattice Energy.
ΔH solution = ΔH hydration (anion + cation) – ΔH lattice
More soluble in water : hydration enthalpy more exothermic than lattice energy.
Less soluble in water : heat of solution more endothermic (positive final result, difficult to dissolve in water).
Factors affect hydration enthalpy and lattice energy: change of ions, size of ions.
Hydration enthalpy and lattice energy both decreases when going down the group.
Change in rates of hydration enthalpy and lattice energy when down group 2:
a. when increase solubility: lattice energy decrease faster than hydration enthalpy.
b. when decrease solubility: hydration enthalpy decrease faster than lattice energy.
When down group 2, the solubility will ____.
Sulphates, SO42-: decreases.
Halide (Cl-, Br-, I-): decreases.
Alkaline Earth Metal: increases (reactivity increases, E^o become more negative, increase reducing agent).
Carbonates, CO3^2-: increases.
Halide (fluoride, F-) : increases.
Hydroxides, OH-: increases.
Oxides O2- : increases. BeO and MgO possess high lattice energy and thus insoluble in water. CaO < SrO < BaO dissolve in water to form basic hydroxides and evolve a large amount of heat.
3. Thermal Stability when down the group 2.
Carbonate, CO32- : increases. charge density of the cations (polarization)
Nitrate, NO3- : increases.
4. Nature of oxides.
When down group 2,
The oxides are basic nature increases from BeO to BaO (due to increasing ionic nature).
BeO (Amphoteric) < MgO (Weak basic) < CaO < SrO < BaO (Strong basic).
BeO dissolves both in acid and alkalies to given salts and is amphoteric.
The oxides of the alkaline earth metals (except BeO and MgO) dissolve in water to form basic hydroxides and evolve a large amount of heat.
BeO and MgO possess high lattice energy and thus insoluble in water.
5. Going down the group, increases atomic radius (size).
Size depends on 2 factors: 1. nuclear charge, 2. screening effect.
Down group 2, the number of protons and the inner electrons increases. Increase in the screening effect more or less (not obvious) cancels out the increase in the nuclear charge. Proton number (nuclear charge) – Screening effect (inner electrons) = Effective nuclear charge.
Be 4-2 = +2
Mg 12-10 = +2
Ca 20-18 = +2
Sr 38-36 = +2
Ba 56-54 = +2
There is little change in the effective nuclear charge.
Attraction between the nucleus and the electron decreases, causing the electron cloud to become far from nucleus causing the atomic radius and ionic radius to increase with increasing proton number.
Why are Group 2 elements Good reducing agents?
The Group 2 elements are powerful reducing agents. All Group 2 elements have 2 electrons in their outer shell. They generally lose these two outershell electrons in order to react and, by doing so, they form M2+ ions.
The values of the standard reduction electrode potential indicate the relative reactivity of the elements:
When going down the group, the size of the atom (atomic radius) increases. The valence electrons are further away from nucleus. Thus, it is easier for the metal to lose electrons (valence electrons to be lost in a reaction). Hence, the metals become increasingly electropositive and their reducing strength increases. Larger the shielding effect, the weaker is the nuclear pull towards the electrons.
More negative value of SREP, more electropositive of the metal.
Group 2 elements have lower standard reducing electrode potential compared to Group 1 elements in the same period. Metal G2 less reactive than G1. 2 electrons must removed when reacts. AE activation energy of the reaction is higher.
Why does beryllium form covalent bonds?
Be2+ ion has small size and high charge. Hence, it has high polarizing power and forms covalent compounds. Other alkaline earth metals form ionic compounds.
Why only beryllium can form complex ions? Why it is not related to the standard reduction electrode potential as the explanation?
Complex ion formation: Definition: A complex ion is a cation bonded to small molecules called ligands by dative bonds. A dative bond is covalent bond in which both the shared electrons are provided by one atom only. The tendency to form complexes by group 2 elements decreases as the size of M2+ ions increases due to the decrease in charge density. For instance, Be2+ on account of its small size and high charge density, forms stable complexes, e.g. [BeF4]^2- while Ba2+ forms very few.
The ability of beryllium to form complex ions does not depend on the value of the standard reduction electrode potential. Beryllium has quite a high electronegativity compared with the rest of the Group. That means that it attracts a bonding pair of electrons towards itself more strongly than magnesium and the rest do. In order for an ionic bond to form, the beryllium has to let go of its electrons. It is too electronegative to do that.
Group 2 alkaline metal have a negative standard reduction electrode potential because:
a. have tendency to lose electrons.
b. work as negative terminal when connected to SHE.