Graph showing the trends of the atomic numbers as you do down the group.
As you go down the group the Atomic Number increases because the number of protons in the nucleus increases.
Evidence of using an advanced internet search
Atomic Radius
Atoms don’t have a definite surface. An atom’s electrons are constantly moving around the nucleus. The atomic radius is one half of the distance between the nuclei of two identical atoms. The unit of measurement is a nanometer. Below is a graph and table showing the atomic radius as you go down the group.
This graph shows that as you go down the group, the Atomic Radius increases. The reason for this is that there are no more filled shells of electrons.
Electronegativity
This is the power of an atom to withdraw electron density from a covalent bond.
The Electronegativity of an element decreases as you go down Group 2 because the bonding electrons are further away from the nucleus, and therefore shielded from the nuclear charge by more inner shells.
First Ionisation Energy
This is the minimum energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to produce 1 mole of gaseous positive ions.
The equation for this is:
(X being the element)
X (g) X+ (g) + e-
Certain factors affect the First Ionisation Energy of an element. These are:
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The charge on the nucleus – the more protons there are in the nucleus the more positively charged the nucleus, and the more strongly electrons are attracted to it.
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The distance of the electron from the nucleus - An electron close to the nucleus will be much more strongly attracted than one further away.
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The number of electrons between the outer electrons and the nucleus – outer electrons can be shielded/screened by the inner electrons.
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Whether the electron is on its own orbital or paired with another electron- two electrons on the same orbital repel each other as they are both the same charge. This means they are easier to remove than those that aren’t paired.
Electron configuration of the first 3 metals in Group 2:
As you go down group 2, the First Ionisation Energy, like electronegativity, decreases because the outer electrons are easier to remove as they are in a higher sub shell. This means that it is shielding from the nuclear charge my more inner shells as it is further away from the nucleus.
Melting points of Group 2 Metals
The Group 2 Metals are all metals with metallic bonding, so you expect their melting points to be high. In metallic bonding, metallic cations in a metallic lattice are attracted to delocalised ions. Going down Group 2, the number of delocalised electrons remains the same, the charge on each metal cation stays the same at 2+, but the ionic radius increases so the attraction between the delocalised electrons and the metal cations decreases.
Therefore the melting point generally decreases as you go down. Beryllium and Magnesium have a different solid structure which lowers their melting points
The Solubility of Metal Sulphates
The sulphates become less soluble as you go down the Group.
The Solubility of Metal Hydroxides
The hydroxides become more soluble as you go down the group. Examples of these trends:
Magnesium hydroxide appears to be insoluble in water. However, if you shake it with water, filter it and test the pH of the solution, you find that it is slightly alkaline. This shows that there are more hydroxide ions in the solution than there were in the original water. This proves that some Magnesium hydroxide must have dissolved.
Calcium hydroxide solution is used as "lime water". 1 litre of pure water will dissolve about 1 gram of calcium hydroxide at room temperature.
Barium hydroxide is soluble enough to be able to produce a solution with a concentration of around 0.1 mol dm-3 at room temperature.
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