17.5-17.6Free
Your guide: Miss SanaCounts atoms the way you count rupees: carefully.
Milk of magnesia barely dissolves in water. Epsom salt dissolves easily. Both are magnesium compounds.
Why? It is a story of two energies.
Two energies, two trends, complex ions.
Solubility depends on a balance of two energies.
Enthalpy of hydration is the energy released when gaseous ions are surrounded by water molecules. It is always exothermic, so it is negative.
Small ions with high charge attract water more strongly. Their hydration enthalpy is more negative. It gets less negative down the group.
Lattice enthalpy is the energy needed to break one mole of a solid ionic compound into gaseous ions. It is always endothermic.
If the two ions are of similar size, the forces are stronger and lattice enthalpy is higher. If one is much larger, the forces are weaker.
Enthalpy of solution is the overall energy change when a solid dissolves. A more negative or less positive value usually favours higher solubility.
Mg(OH)2 is sparingly soluble. Ba(OH)2 is quite soluble.
Down the group both the hydration enthalpy and the lattice enthalpy decrease. For hydroxides the decrease in lattice enthalpy is the more significant one.
So it gets easier to break the lattice, and solubility increases.
Table 17.4, in mol per 100 g of water: Mg(OH)2 2 × 10−5, Ca(OH)2 1.5 × 10−3, Sr(OH)2 3.4 × 10−3, Ba(OH)2 1.5 × 10−2.
MgSO4 is soluble. BaSO4 is insoluble.
The sulphate ion is relatively large. Down the group both enthalpies decrease. But for a large anion the decrease in hydration enthalpy is the more significant one.
So less energy is paid back by hydration, and solubility decreases.
Table 17.4, in mol per 100 g of water: MgSO4 1.83 × 10−1, CaSO4 4.66 × 10−3, SrSO4 7.11 × 10−5, BaSO4 9.43 × 10−7.
A complex ion has a central metal ion bonded to several electron-donating species.
The metal ion is a Lewis acid, an electron pair acceptor. The donors, such as H2O, NH3, Cl− and OH−, are Lewis bases.
The number of coordinate bonds is the coordination number. Common values are 2, 4 and 6.
Complexes are more typical of transition metals. Group 2 ions form them less, especially the ones at the top of the group.
Small ions have high charge density, so they attract donors better. All Group 2 ions have a +2 charge.
In water the ions are hydrated. This is a complex ion with water as the donor.
Beryllium is the smallest. It gives [Be(H2O)4]2+, coordination number 4.
Magnesium gives [Mg(H2O)6]2+, coordination number 6. Ca, Sr and Ba mostly give 6 too, but the interaction gets weaker down the group.
EDTA is a donor with six electron pairs. It wraps around a calcium ion in a ring structure.
Ca2+(aq) + EDTA4−(aq) → [Ca(EDTA)]2−(aq)
This is used to soften water. EDTA removes Ca2+ and Mg2+ from hard water. It is also used in medicines to remove metal ions from the body.
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Milk of magnesia barely dissolves in water. Epsom salt dissolves easily. Both are magnesium compounds.
Why? It is a story of two energies.
Two energies, two trends, complex ions.
Solubility depends on a balance of two energies.
Enthalpy of hydration is the energy released when gaseous ions are surrounded by water molecules. It is always exothermic, so it is negative.
Small ions with high charge attract water more strongly. Their hydration enthalpy is more negative. It gets less negative down the group.
Lattice enthalpy is the energy needed to break one mole of a solid ionic compound into gaseous ions. It is always endothermic.
If the two ions are of similar size, the forces are stronger and lattice enthalpy is higher. If one is much larger, the forces are weaker.
Enthalpy of solution is the overall energy change when a solid dissolves. A more negative or less positive value usually favours higher solubility.
Mg(OH)2 is sparingly soluble. Ba(OH)2 is quite soluble.
Down the group both the hydration enthalpy and the lattice enthalpy decrease. For hydroxides the decrease in lattice enthalpy is the more significant one.
So it gets easier to break the lattice, and solubility increases.
Table 17.4, in mol per 100 g of water: Mg(OH)2 2 × 10−5, Ca(OH)2 1.5 × 10−3, Sr(OH)2 3.4 × 10−3, Ba(OH)2 1.5 × 10−2.
MgSO4 is soluble. BaSO4 is insoluble.
The sulphate ion is relatively large. Down the group both enthalpies decrease. But for a large anion the decrease in hydration enthalpy is the more significant one.
So less energy is paid back by hydration, and solubility decreases.
Table 17.4, in mol per 100 g of water: MgSO4 1.83 × 10−1, CaSO4 4.66 × 10−3, SrSO4 7.11 × 10−5, BaSO4 9.43 × 10−7.
A complex ion has a central metal ion bonded to several electron-donating species.
The metal ion is a Lewis acid, an electron pair acceptor. The donors, such as H2O, NH3, Cl− and OH−, are Lewis bases.
The number of coordinate bonds is the coordination number. Common values are 2, 4 and 6.
Complexes are more typical of transition metals. Group 2 ions form them less, especially the ones at the top of the group.
Small ions have high charge density, so they attract donors better. All Group 2 ions have a +2 charge.
In water the ions are hydrated. This is a complex ion with water as the donor.
Beryllium is the smallest. It gives [Be(H2O)4]2+, coordination number 4.
Magnesium gives [Mg(H2O)6]2+, coordination number 6. Ca, Sr and Ba mostly give 6 too, but the interaction gets weaker down the group.
EDTA is a donor with six electron pairs. It wraps around a calcium ion in a ring structure.
Ca2+(aq) + EDTA4−(aq) → [Ca(EDTA)]2−(aq)
This is used to soften water. EDTA removes Ca2+ and Mg2+ from hard water. It is also used in medicines to remove metal ions from the body.
Q1. Quick Check 17.3 (a): why does MgSO4 dissolve more easily than BaSO4?
Sulphate solubility decreases down the group. The sulphate ion is large. Down the group the hydration enthalpy falls more than the lattice enthalpy. So BaSO4 gains less from hydration and is insoluble. MgSO4 is 1.83 × 10−1, BaSO4 is 9.43 × 10−7.
Q2. Quick Check 17.3 (c): which is more soluble, Ca(OH)2 or Sr(OH)2? Give the reason.
Sr(OH)2: 3.4 × 10−3 against 1.5 × 10−3 for Ca(OH)2. Hydroxide solubility increases down the group. The lattice enthalpy falls more than the hydration enthalpy.
Q3. Exercise Q2 (b): dichromate is larger than sulphate. Which is more soluble, calcium dichromate or strontium dichromate?
Follow the sulphate pattern, because dichromate is also a large anion. Solubility decreases down the group. So calcium dichromate is more soluble. (Answer built from the book's rule. The book prints no key.)
Q4. Exercise Q2 (a): 0.1 mol each of Ca(OH)2 and Ba(OH)2 in separate 100 cm3 of water. Which solution has the higher pH?
Ba(OH)2. It is more soluble (1.5 × 10−2 against 1.5 × 10−3 mol per 100 g), so more OH− ions are in solution. (Reasoning from Table 17.4. The book prints no key.)
✗ “Hydroxides and sulphates follow the same solubility trend.”
✓ They are opposite. Hydroxides increase down the group. Sulphates decrease.
✗ “Hydration enthalpy is positive.”
✓ It is always exothermic, so it is negative.
✗ “Only transition metals form complex ions.”
✓ Group 2 ions form them too, for example [Mg(H2O)6]2+. It is less pronounced than in transition metals.
1. Which sulphate is the least soluble?
(a) Table 17.4: BaSO4 is 9.43 × 10−7.
2. Hydroxide solubility down Group 2:
(a) From 2 × 10−5 to 1.5 × 10−2.
3. Why does Be2+ form complexes more easily than Ba2+?
(a) Small size and the same charge give a high charge to size ratio.
4. The coordination number of [Mg(H2O)6]2+ is:
(a) Six water molecules are bonded.
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