Trend In Atomic Radius

Going down Group 2:

There are more filled energy levels between the nucleus and the outer electrons, therefore the outer electrons are more shielded from the attraction of the nucleus So the electrons in the outer energy levels are further from the nucleus and the atomic radius increases.

As the number of protons in the nucleus increases going down Group 2, you might expect the atomic radius to decrease because the nuclear charge increases.

This does not happen, because although the electrons in the inner energy levels become closer to the nucleus, the factors described above have a greater influence on the atomic radius overall.

Reaction with Air

Magnesium burns vigorously with a bright white flame when strongly heated in air/oxygen to form a white powder of magnesium oxide.

Magnesium + oxygen magnesium oxide

2Mg(s) + O₂(g) 2MgO(s) Calcium burns quite fast with a brick red flame when strongly heated in air/oxygen to form the white powder calcium oxide.

Calcium + oxygen calcium oxide

2Ca(s) + O₂(g) 2CaO(s)

Reaction with Water

1. Magnesium will not react with cold water. Even finely powdered magnesium reacts only very slowly. Magnesium will react with gaseous water (steam) to form magnesium oxide and hydrogen. Magnesium + steam magnesium oxide + hydrogen.

Mg(s) + H₂O(g) MgO(s) + H₂(g)

Magnesium oxide is a base. It will not dissolve in water.

In fact magnesium is so reactive; it will even burn in carbon dioxide, the products being white magnesium oxide powder and black specks of elemental carbon!

Magnésium + carbon dioxide ==> magnesium oxide + carbon 2Mg(s) + CO₂(g) ==> 2MgO(s) + C(s)

2. Calcium (and the metals below calcium in group 2) will react with cold water. They will sink as they react, unlike the group 1 metals which float.

Calcium + water calcium hydroxide + hydrogen. Ca(s) + 2H₂O(l) Ca(OH)₂(s) + H₂(g)

Calcium hydroxide is called slaked lime and will dissolve a little in water to form lime water

Reaction with Acids

Magnesium is very reactive with dilute hydrochloric acid forming the colourless soluble salt magnesium chloride and hydrogen gas.

Magnesium + hydrochloric acid ==> magnesium chloride + hydrogen Mg(s) + 2HCl(aq) ==> MgCl₂(aq) + H₂(g)

Magnesium is very reactive with dilute hydrochloric acid forming the colourless soluble salt calcium chloride and hydrogen gas.

Calcium + hydrochloric acid ==> calcium chloride + hydrogen

Ca(s) + 2HCl(aq) ==> CaCl2(aq) + H2(g)

Not very reactive with dilute sulphuric acid because the colourless calcium sulphate formed is not very soluble and coats the metal inhibiting the reaction, so not many bubbles of hydrogen.

Calcium + sulphuric acid ==> calcium sulphate + hydrogen

Ca(s) + H₂SO₄(aq) ==> CaSO₄(aq/s) + H₂(g)

Reaction with Halogens

They occur in nature only in compounds because of their high reactivity.

They are less reactive than group 1 elements due to higher IE.

They react with elements in group 7 to give the general formula MX₂ (M is the metal and X represents any members of group 7.

M + X₂ –> MX₂ Example Mg + Br₂ –> MgBr₂

Summary of the Reactivity Trend of Alkali Earth Metals

1. All Group II metals have 2 valence electrons (2 electrons in the highest energy level)
2. Atomic radius increases down the Group as successive ‘electron shells’ (energy levels) are filled
3. Down the Group, first ionization energy (the energy required to remove 1 electron from the gaseous atom) decreases.
4. As the atomic radius increases and the electron is further from the nucleus it is less attracted to the nucleus (electron is said to be ‘shielded’)
5. Down the Group, second ionization energy (the energy required to remove an electron from the gaseous positive ion) decreases.

As successive electron ‘shells’ (energy levels) are filled, the electron is further from the positively charged nucleus, and therefore less attracted to it, making the electron easier to remove.

1. Down the Group, third ionization decreases (the energy required to remove an electron from the gaseous ion of charge 2+).
2. As successive electron ‘shells’ (energy levels) are filled, the electron is further from the positively charged nucleus, and therefore less attracted to it, making the electron easier to remove.
3. Second ionization energy is higher than the first ionization for each element.
4. This is because it is harder to remove the electron since there are more positive charges (protons) in the nucleus than there are negative charges (electrons in ‘shells’), hence the electron’s attraction to the nucleus is greater.
5. Third ionization energy is substantially higher than the second ionization energy.

When 2 electrons have been removed from the gaseous atom, the remaining electrons are arranged like a noble gas, which is a very stable electron configuration.

1. It is very difficult to remove an electron from this arrangement.
2. In general, electro negativity decreases down the Group as successive energy levels (electron shells) are filled, the positive nucleus exerts less force on electrons and so has less ability to attract electrons.
3. Melting point decreases down the Group as the elements become less metallic in nature.]