As you go down the group from one element down to the next the atomic radius gets bigger due to an extra filled electron shell. The outer electron is further and further from the nucleus and is also shielded by the extra full shell of negative charge, therefore the outer electron is less and less strongly held by the positive nucleus as the attractive force is decreased, and so.
Table of Contents
This combination of factors means the outer electron is more easily lost, the M+ ion more easily formed, and so the element is more reactive as you go down the group
Uses of Alkali Metals
Sodium Na+ salts
- Common salt from sea water or underground deposits is sodium chloride and is the raw material for making sodium, hydrogen, chlorine and sodium chloride by electrolysis.
- Sodium hydrogen carbonate (NaHCO3) Used in baking soda, pharmaceutical products like indigestion tablets and fire extinguishers.
- Sodium hydroxide (NaOH) Used in the manufacture of soaps, detergents, salts of acids, paper and ceramics.
The Alkaline Earth Metals – Group 2 – Properties.
The second group (group II A) has Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra). They have two electrons in their last shell and their valence is +2 as they give up two electrons to form compounds. The elements in group II A are not as metallic as the alkali metals. They form oxides easily and are known as alkali earth metals.
Group 2 Alkaline Earth Metals
The alkaline earth metals are all reactive, losing their two outer electrons to form a 2+ ion with non-metals.
Mg Mg2+ + 2e-
Ca Ca2+ + 2e-
From calcium going down the group, they have to be stored under oil, or they react with oxygen in the air. They are less reactive than the alkali metals (Group 1). Calcium and magnesium are fourth and fifth in the reactivity series.
They all have the common properties of metals, being silvery-grey in colour, and good conductors of heat and electricity. They are less soft than the alkali metals, and it is difficult to cut them with a knife. Strontium, barium and radium are all too reactive or unstable to be used. All you need to know about these three is that they have the same chemical properties as magnesium and calcium
It would be expected to lose its two outer electrons like the rest of Group 2 but beryllium is so small that it doesn’t like to lose two electrons.
Its compounds have covalent character! As we proceed to group III and further, we will notice that the number of valence electrons increases by one in each subsequent group.
Explanation of this trend
The first ionization energy is the enthalpy change when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge. It is an endothermic process, i.e. is positive.
A general equation for this enthalpy change is:
Ionization energy is governed by:
- The charge on the nucleus,
- The amount of screening by the inner electrons,
- The distance between the outer electrons and the nucleus.
Going down Group 2:
How does the first ionization energy change going down the group? The outer electrons are held in their shells by the attractive force of the positive protons in the nucleus, the nuclear attraction.
As more and more electron shells are added this force gets weaker because
- The distance between the outer electrons and the nucleus is increasing
- The inner electrons shield the nuclear electrons from the outer electrons, electronic shielding.
The lower the ionization energy the easier it is to remove electrons from the outermost shell of the atom. As you go down a group the ionization energy decreases. This also explains why metals get more reactive as you go down a group. It gets easier for them to give up electrons to form bonds.
As the number of protons in the nucleus increases going down Group 2, you might expect the first ionization energy to increase because the nuclear charge increases. This does not happen, because the factors described above have a greater influence on the value of the first ionization energy.