Periodic Table: Trends in Group 2 Elements (alkaline earth metals) Chemistry Tutorial

Key Concepts

• The vertical columns in the periodic table of the elements are known as groups.
• The second vertical column from the left in the periodic table is referred to as Group 2.

  Alternative names for Group 2 are:
  (a) alkali earth metals (still commonly used)
  (b) alkaline earth metals, or alkaline-earth metals ⁽¹⁾ (still commonly used)
  (c) group IIA (this numbering is no longer used)

• The name and symbol for the elements in Group 2 are given below:

  Group 2 Elements (Alkaline-earth Metals)
  name	symbol
  beryllium	Be
  magnesium	Mg
  calcium	Ca
  strontium	Sr
  barium	Ba
  radium	Ra

• Except for beryllium⁽²⁾, the Group 2 elements are typical metals:

  (a) relatively soft, but harder than group 1 metals, shiny solids at room temperature and pressure that are good conductors of heat and electricity

  (b) Moderately-high melting point.

  (c) have 2 valence electrons (2 electrons in the highest energy level)

  (d) are very reactive

  (e) form cations with a charge of +2 (M²⁺) when they combine with non-metals in an ionic compound

  (f) form white ionic compounds ⁽³⁾

• Going down group 2 from top to bottom the elements display the following general trends:

  (a) Atomic radius increases.

  (b) First ionization energy decreases.

  (c) Second ionization energy decreases.

  Second ionization energy is about double the first ionization energy for each element.

  (d) Third ionization energy decreases.

  Third ionization energy is much greater, about 4 times greater, than the second ionization energy for each element.

  (e) Electronegativities decrease as successive energy levels (electron shells) are filled resulting in the positive nucleus exerting less of a force of attraction on electrons.

  (f) Chemical reactivity increases

  (g) Densities generally increase.⁽⁴⁾

• Summary of trends in Group 2 elements:

  Trends: Decreasing				Group 2 Elements		Trends: Increasing
  First Ionisation Energy	Second Ionisation Energy	Third Ionisation Energy	Electro- negativity	name	symbol	atomic number	atomic radius	Reactivity
  largest	largest	largest	highest	beryllium	Be	lowest	smallest	less reactive
  ↑	↑	↑	↑	magnesium	Mg	↓	↓	↓
  ↑	↑	↑	↑	calcium	Ca	↓	↓	↓
  ↑	↑	↑	↑	strontium	Sr	↓	↓	↓
  smallest	smallest	smallest	lowest	barium	Ba	highest	largest	most reactive

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Table of Data for Group 2 Elements

The table below gives the name, atomic number, electronic configuration of the atom, the first, second and third ionisation energy, melting point, density and electronegativity, of the Group 2 elements (alkaline-earth metals).

Carefully inspect this data to find trends, or patterns, in the properties of group 2 elements.

These patterns, or trends, recur throughout the periodic table and are referred to more generally as periodic trends, or, as periodicity.

Trends in Electronic Configuration of Group 2 Elements

Consider the electronic configuration of group 2 elements. Can you see a trend (a pattern)?

Atoms of group 2 elements have just 2 electrons in the highest energy level (also known as the valence shell of electrons).

It is even easier to see this if we use a short-hand description of the electronic configuration of each atom in which the electrons that make up part of a Noble Gas (group 18) electron configuration are represented in square brackets followed by the number of electrons in the valence shell.
We have done this in the table below:

If an atom (M) of a group 2 element lost both these valence electrons (2e^-), then the ion of the group 2 element would have a charge of +2 (M²⁺) as shown in the equations below:

And, the positively charged ion (cation) formed would have the same electronic configuration as a group 18 (Noble Gas) element, we say that the cation is isoelectronic with the Noble Gas, as shown below:

and the cation of a group 2 element would therefore be chemically very stable (that is, no longer very reactive), just like a Noble Gas (group 18 element).

So, just how likely is it that a group 2 element will lose both valence electrons and form a cation .....

Trends in Ionisation Energy of Group 2 Elements

Ionisation energy (or ionization energy) is the energy required to remove an electron from a gaseous species.

First ionisation energy (or first ionization energy) refers to the energy required to remove an electron from a gaseous atom.

We can write a general equation to describe the removal of an electron (e^-) from a gaseous atom (M_(g)) to produce a gaseous cation with a charge of +1 (M⁺_(g)) as:

M_(g) → M⁺_(g) + e^-

Second ionisation energy refers to the energy required to remove an electron (e^-) from the gaseous ion with a charge of +1 (M⁺_(g)) to form a gaseous ion with a charge of +2 (M²⁺_(g)) as shown in the equation below:

M⁺_(g) → M²⁺_(g) + e^-

If the value of the ionisation energy is high, then lots of energy is required to remove the electron, and the reaction is less likely to occur readily.
If the value of the ionisation energy is low, then little energy is required to remove the electron, and the reaction is more likely to occur readily.

So let's look at the values of the first and second ionisation energy for each Group 2 element (alkaline-earth metal):

As you go down group 2 from top to bottom, the value of first ionisation energy decreases, it is progressively easier to remove the first valence electron.

As you go down group 2 from top to bottom, the value of the second ionisation energy decreases, it is progressively easier to remove the second valence electron.

The suggestion here is that the chemical reactivity of the elements increase as you go down group 2 from top to bottom.
That is, since it requires less energy to remove the two valence electrons as you go down the group, the chemical activity of these elements will increase going down the group.

You might also notice that the value of the second ionisation energy for each element is about double that of the first ionisation energy.

If we are right and the electronic configuration of a Noble gas (Group 18) element is particularly stable, then it should be very difficult, that is, require a lot more energy, to remove the third electron from each Group 2 element.

Third ionisation: M²⁺_(g) → M³⁺_(g) + e^-

So, let's look at the value of each third ionization for each group 2 element:

In general, it requires a bit less than twice as much energy to remove the second valence electron than it does to remove the first valence electron from a gaseous atom of each element.

But in general it requires more than double this amount of energy again in order to remove the third electron.
This strongly supports the concept that the electronic configuration of a Noble Gas (group 18) element is remarkably stable and that any atom or ion with this structure will not be chemically reactive.

As a result, Group 2 elements form ionic compounds in which the group 2 cation has a charge of 2+. ⁽⁵⁾

But why is it easier to remove these valence electrons as you go down group 2 from top to bottom....

Trends in Atomic Radius of Group 2 Elements

First, lets think about the number of electron shells (or energy levels) being filled to make an atom of each group 2 element:

As you go down group 2 from top to bottom, you are adding a whole new "electron shell" to the electronic configuration of each atom.
Surely that will increase the size of each atom as you go down the group?
We record the "size" of an atom using its "atomic radius".
Consider the values for the atomic radius of each of the atoms in group 2 as shown in the table below:

As you go down group 2 from top to bottom the radius of the atom of each successive element increases.
This means that the negatively charged valence electrons get further away from the positively charged nucleus and we say that these electron are 'shielded'.
So, the positively charged nucleus has less of a "pull" on the valence electrons as you go down the group.
Therefore, the valence electrons are easier to remove, and therefore the ionisation energy decreases down the group as discussed in the previous section.

All of this means that the reactivity of Group 2 elements increases as you go down the group from top to bottom...

Trends in Reactivity of Group 2 Elements (alkaline-earth metals)

All the group 2 elements (M_(s)), except beryllium, react with water (H₂O_(l)) to form hydrogen gas (H_2(g)) and an alkaline (basic) aqueous solution (M(OH)_2(aq)) as shown in the balanced chemical equations below:

The reaction between magnesium and water is usually slow because magnesium readily reacts with oxygen and a protective layer of magnesium oxide forms over the metal.
The reactions between other Group 2 elements and water is vigorous.

Beryllium and magnesium do not combine directly with hydrogen, however, calcium, strontium and barium will combine directly with hydrogen:

Reactions with water and hydrogen as described above indicate that there is a general trend in the chemical reactivity of group 2 elements: the reactivity of the group 2 elements increases as you go down the group from top to bottom.

The group 2 metals (M_(s)) react with oxygen gas (O_2(g)) at room temperature and pressure to form oxides with the general formula MO as shown in the balanced chemical reactions below:

Group 2 metals (M_(s)) react with halogens (group 17 elements) to form halides with the formula MX₂.
For example, group 2 elements react with the halogen chlorine gas (Cl_2(g)) to form an ionic chloride⁽⁶⁾ (MCl_2(s)) as shown in the balanced chemical equations below:

Group 2 elements will also combine with sulfur to form sulfides with the general formula MS:

and they will combine with nitrogen to form nitrides with the general formula M₃N₂:

Worked Example 2

Problem Solving using StoPGoPS method