Now let's look at $\ce{SO4^2-}$. Nitrates of group -1 and group-2 metals are all soluble in water. Charge Density and Polarising Power of Group 2 Metal Cations. Due to this, the solubility increases with increase in the molecular weight on moving down the group. This is why the solubility of Group 2 hydroxides increases while progressing down the group. Hence, more is the stability of oxide formed, less will be stability of carbonates. Place 2 g of a Group 2 metal carbonate in the test tube and reweigh. The hydroxides of the Group II metals, which may be used in thermochemical water-splitting cycles, have been investigated thermoanalytically. The solubilities of these salts further increase on descending the group. Looking at the enthalpy change of formation for group 2 metal oxides it’s clearly less energy is needed to break them as you go down the group. Thus stability of alkaline earth metal hydroxides decreases with decrease in lattice enthalpy as the size of alkali earth metal cations increases down the group. A higher temperature is required to decompose Ba(NO 3) 2 as compared to Mg(NO 3) 2. So what causes this trend? Weigh a test tube. The thermal stability of the hydrogencarbonates. ... Solubility of the carbonates increases as you go down Group 1. Even for hydroxides we have the same observations. Correct option: (d) Ba(OH) 2 < Sr(OH) 2 < Ca(OH) 2 < Mg(OH) 2 Explanation: Stability of ionic compounds decreases with decrease in lattice enthalpy. The size of B e 2 + is smallest and the size of B a 2 + is highest. The solubility of alkaline metal carbonates and sulphates decreases with decrease in hydration energy as we move down the group. Decomposition temperatures and decomposition enthalpies of the four hydroxides increase with increasing atomic weight of the compounds. The least soluble hydroxide in Group 1 is lithium hydroxide - but it is still possible to make a solution with a concentration of 12.8 g per 100 g of water at 20°C. BeCO 3 ⇌ BeO + CO 2. Thus stability of alkaline earth metal hydroxides decreases with decrease in lattice enthalpy as the size of alkali earth metal cations increases down the group. 2 M N O 3 h e a t 2 M n O 2 + O 2 This page looks at the solubility in water of the hydroxides, sulphates and carbonates of the Group 2 elements - beryllium, magnesium, calcium, strontium and barium. Although it describes the trends, there isn't any attempt to explain them on this page - for reasons discussed later. Solution: Stability of ionic compounds decreases with decrease in lattice enthalpy. (ii) Thermal stability Alkali and alkaline earth metal nitrates decompose on heating. The respective TG- and DSC-curves are represented. Attach the delivery tube to the test tube. Magnesium hydroxide: this is the most insoluble and can be brought as a suspension in water. There is no reaction or precipitate when dilute sodium hydroxide is added to a solution of Sr 2+ or Ba 2+ ions. The same thing applies to the cation while progressing down the group. Place the other end of the delivery tube into a test tube which is one third full of limewater. Sulphates: Thermal stability The sulphates of group-1 and group-2 metals are all thermally stable. Let's use MgCO 3 as an example. —————————————————— Uses of sulphate and hydroxides. Alternative Thermal decomposition of group 2 carbonates practical. Hence, barium hydroxide is more soluble than beryllium hydroxide. As the size increases, the decrease in the lattice energy is much more than the decrease in the hydration energy. On heating alkali metal (Na, K, Rb and Cs) decompose to form metal nitrites and oxygen. The hydroxides. solubility: sulphates of alkali metals are soluble in water. 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