Periodic Table: Trends in Properties of Group 1 Elements (alkali metals) Chemistry Tutorial

Key Concepts

• The vertical columns in the periodic table of the elements are known as groups.
• The first vertical column in the periodic table is referred to as Group 1.

  Alternative names for Group 1 are:
  (a) alkali metals⁽¹⁾ (still commonly used)
  (b) group IA (this numbering is no longer used)

• The name and symbol for the elements in Group 1 are given below:

  Group 1 Elements (Alkali Metals)
  name	symbol
  lithium	Li
  sodium	Na
  potassium	K
  rubidium	Rb
  caesium(2)	Cs
  francium(3)	Fr

• All Group 1 elements:

  (a) are soft, solid, shiny metals at room temperature and pressure that are good conductors of heat and electricity

  (b) have 1 valence electron (1 electron in the highest energy level)

  (c) are very reactive

  (d) form cations with a charge of +1 (M⁺) when they combine with non-metals in an ionic compound

  (e) form white ionic compounds ⁽⁴⁾

• Going down group 1 from top to bottom the elements display the following trends:

  (a) Atomic radius increases.

  (b) First ionization energy decreases.

  (c) Second ionization energy decreases.

  Second ionization energy is much greater than the first ionization energy for each element.

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

  (e) Chemical reactivity increases

  (f) Melting point decreases.

  (g) Densities generally increase.

• Summary of trends in Group 1 elements:

  Trends: Decreasing			Group 1 Elements		Trends: Increasing
  Melting Point	First Ionisation Energy	Electronegativity	name	symbol	atomic number	atomic radius	Reactivity
  highest	largest	highest	lithium	Li	lowest	smallest	less reactive
  ↑	↑	↑	sodium	Na	↓	↓	↓
  ↑	↑	↑	potassium	K	↓	↓	↓
  ↑	↑	↑	rubidium	Rb	↓	↓	↓
  ↑	↑	↑	caesium	Cs	↓	↓	↓
  lowest	smallest	lowest	francium	Fr	highest	largest	most reactive

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Table of Data for Group 1 Elements

The table below gives the name, atomic number, electronic configuration of the atom, the first and second ionisation energy, melting point, density and electronegativity, of the Group 1 elements (alkali metals).

Carefully inspect this data to find trends, or patterns, in the properties of group 1 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 1 Elements

Consider the electronic configuration of group 1 elements. Can you see a trend (a pattern)?

Atoms of group 1 elements have just 1 electron 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 1 element lost that valence electron (e^-), then the ion of the group 1 element would have a charge of +1 (M⁺) 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 1 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 1 element will lose that valence electron and form a cation .....

Trends in Ionisation Energy of Group 1 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^-

So, the first ionisation energy for lithium refers to the energy required to remove 1 electron (e^-) from an atom of lithium which is in the gaseous state (Li_(g)).
The products of the reaction are an electron and a gaseous lithium ion with a charge of +1 (Li⁺_(g)).

We can represent the first ionisation of each group 1 element as shown below:

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

So, let's look at the values for the first ionisation energy for each group 1 element:

As you go down group 1 from top to bottom, it gets easier to remove the valence electron and form the positively charged cation.
Group 1 elements increase in chemical reactivity as you go down the group from top to bottom.

We have evidence for the stability of the electronic configuration of the group 1 cations based on inspection of the values for the second ionisation for this group.
Second ionisation energy refers to the amount of energy required to remove an electron (e^-) from the gaseous singly charged cation (M⁺_(g)) to form a gaseous cation with a charge of 2+ (M²⁺_(g)):

M⁺_(g) → M²⁺_(g) + e^-

Let's compare the values for the first ionisation energy and the second ionisation energy for each element in group 1:

Note that second ionisation decreases down the group, just like first ionisation energy, but, the values for the second ionisation energy are much, much, larger than the values for the first ionisation energy.
It is about 10 times harder to remove an electron from the M⁺_(g) ion compared to removing an electron from the M_(g) which provides evidence for the stability of the electron configuration of the M⁺_(g) ion.

But why is that 1 valence electron easier to remove as you go down group 1 .....

Trends in Atomic Radius of Group 1 Elements

First, lets think about the number of electron shells (or energy levels) being filled to make an atom of each group 1 element:

As you go down group 1 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 1 as shown in the table below:

As you go down group 1 from top to bottom the radius of the atom of each successive element increases.
This means that the negatively charged valence electron gets further away from the positively charged nucleus and w say that the electron is 'shielded'.
So, the positively charged nucleus has less of a "pull" on the valence electron as you go down the group.
Therefore, the valence electron is easier to remove, and therefore the ionisation energy decreases down the group as discussed in the previous section.

All of this makes Group 1 metals very reactive..... but just how reactive are they?

Trends in Reactivity of Group 1 Metals

All Group 1 metals react with water (if you haven't seen this then you should go search for some YouTube videos).
We can represent the overall reaction of a group 1 metal (M_(s)) with water (H₂O_(l)) to form an aqueous metal hydroxide (MOH_(aq)) and hydrogen gas (H_2(g)) as:

If you cut off a thin slice of lithium and place it in a beaker of room temperature water the reaction will take place slowly, you will see bubbles of hydrogen gas being produced.

Cut off a thin slice of sodium and place it in room temperature water and the piece of sodium will whiz around the water because the reaction producing the hydrogen gas is a bit more vigorous.

If you do the same with a thin fresh slice of potassium the reaction is even more vigorous, it will probably produce a flame, maybe an audible "pop".

If you do the same thing with a thin fresh slice of caesium it will definitely "pop" and produce flame!
The pop is the explosion due to the rapid production, and ignition, of hydrogen gas!

This is a demonstration to show that the reactivity of group 1 metals with water increases as you go down the group from top to bottom.

The results are summarised in the table below:

Group 1 metals (alkali metals) will react with lots of non-metals, even oxygen (O_2(g)) in the atmosphere as shown below:

The Group 1 metals (alkali metals) react so readily with water and oxygen in the atmosphere that storage of these elements is a problem!
This is why group 1 elements are stored in jars filled with a "water-hating"⁽⁷⁾ hydrocarbon solvent such as paraffin oil, cyclohexane or kerosene.⁽⁸⁾

Trends in the Density of Group 1 Elements

Density refers to how much mass of substance is present in a given volume.
Density of a solid is usually measured in units of grams per cubic centimetre (g cm^-3).
Consider the density of group 1 elements as given in the table below:

If we took a cube of lithium measuring 1 cm × 1 cm × 1 cm, then this cube would have a mass of 0.54 g.
A 1 cm × 1 cm × 1 cm cube of sodium would have a greater mass, 0.97 g

As you go down group 1 from top to bottom, the mass of a cubic centimetre of element has a tendency to increase.
That is, the density of group 1 elements shows a "general trend" of increasing as you go down the group from top to bottom.

As you go down group 1 from top to bottom, the mass of the element present per unit volume, in general, increases.
It should be noted that the density of group 1 (alkali metals) is less than that of transition metals because of the group 1 elements' larger atomic radii.
For example, the density of iron, a transition metal, is about 7.87 g cm^-1.

Trends in the Melting Point of Group 1 Elements

At 25°C and normal atmospheric pressure (100 kPa), group 1 metals exist as solids.
The atoms of each element occupy a place within a 3-dimensional array, or metallic lattice, of atoms.
The atoms of metals are held together in the lattice by metallic bonds.
If enough heat energy is supplied to discrupt this arrangement of atoms, the regularity of the lattice breaks down and the solid metal melts.
The melting point of a metal therefore indicates how much energy needs to be supplied to melt the solid metal.
A high melting point means lots of energy is required to melt the solid, but a low melting point means little energy is required to melt the solid.
We can then infer that the interactions between the metal atoms in a high melting point solid must be greater than the interactions between atoms in low melting point solid.
So a high melting point suggests the metallic bonds between metal atoms is stronger, while a lower melting point suggest the metallic bonds between the metal atoms are weaker.

So, let's compare the melting points of our group 1 metals..

First of all we would note that none of the melting points are very high compared to other metals, for example, the melting point of iron is about 1500°C!
This is because Group 1 metals have only 1 electron to contribute the delocalised "sea of electrons" making up the metallic bond and because group 1 metal atoms tend to be larger than other metal atoms it means that these delocalised electrons are further away from the nucleus, so the metallic bond of Group 1 metals is generally weaker than of other metals.

We can identify a trend in the melting points of group 1 elements: the melting point decreases as you go down the group from top to bottom.
As the atomic radius increases down the group, the delocalised electrons making up the metallic bond get further from the nucleus so the metallic bond gets weaker and easier to weaken as you go down the group.

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