# Group 1: Alkali metals

### Group 1: Alkali metals

#### Lessons

In this lesson, we will learn:
• To recall properties of the alkali metals
• To understand the trend in properties found in the alkali metals
• To apply knowledge of electronic structure and bonding to explain the trends in group 1 properties.
• To know the common reactions of group 1 and group 2 elements.

Notes:
• We saw earlier that the Periodic Table is arranged, left to right, by proton number and number of outer shell electrons. The number of outer shell electrons dictates the chemical properties of an element.
Therefore, it is easy to see which elements have similar properties to each other – they will be in the same column of the table as each other, the columns which we call groups.
Groups 1 and 2 have 1 and 2 electrons in the outer shell, respectively. Their behaviour in chemical reactions is similar because of the similar outer shell configuration.

• The alkali metals in group 1 and the alkaline earth metals in group 2 are very well-studied groups of elements, with clear patterns in how their properties change.
They have been known for quite a long time through their compounds, like their metal oxides, because they are reactive and quickly form compounds with oxygen in air. Their oxides all produce alkaline solutions in water, which is how they get their names. Because of their reactivity, they were not isolated as metals until later on.

• Alkali metals and alkaline earth metals have the following properties:
• They are relatively soft metals.
• They are relatively low density metals.
• They have relatively low melting points compared to metals in general.
• They are reactive, more so than d-block metals in general and they react vigorously with water.
• The products of this reaction are hydrogen gas and a metal hydroxide – this forms an alkaline solution, which gives the two groups their names.

• As you go down the group, the properties of the elements change in the following ways:
• The metals get softer.
• The melting point of the metals gets lower.
• The metals get denser.
• The metals get more reactive.
• Ionization energies of the elements decrease.

• The alkali metals all have a valence of 1 and alkaline earth metals have a valence of 2. As they are metals, they form ionic compounds with non-metals. In these compounds, you’ll see alkali metals with a 1+ charge (such as in NaCl), whereas alkaline earth metals will hold a 2+ charge (such as in MgCl2).

• As you descend the table, a similar trend in the change in properties is observed in both groups. For example, both group 1 and group 2 show a decrease in ionization energies going down the group.

As explained in Periodic Trends: Ionization Energy, this is due to greater shielding from more inner electron shells between the nucleus and the outer shell. Greater shielding makes losing the one (in group 1) or two (in group 2) outer shell electrons increasingly easy, and therefore reactivity in general increases going down the two groups.

• The reactions of group 1 and group 2 metals are very similar because they have similar outer shell configurations and elements in both groups are prone to losing their outer shell electrons.
• One of the most common reactions with group 1 and 2 elements is the reaction with water. Group 1 and group 2 metals are both known for reacting vigorously with water compared to other metals. The equation can be written generally as:

For a group 1 metal:
2M (s) + 2H2O (l) $\,$$\,$ 2MOH (aq) + H2 (g)

For a group 2 metal:
M (s) + 2H2O (l) $\,$$\,$ M(OH)2 (aq) + H2

The stoichiometry is slightly different because two hydroxide groups will be bonded to the group 2 metal because it forms a 2+ ion, unlike the 1+ ion of a group 1 metal.

• Another common reaction is the group 1 or 2 reaction with oxygen. This can be written generally as:

For a group 1 metal:
4M (s) + O2 (g) $\,$$\,$ 2M2O (s)

For a group 2 metal:
2M (s) + O2 (g) $\,$$\,$ 2M2O (s)

• Again, the stoichiometry is slightly different in the two reactions.
In the reaction with heavier group two metals, a metal peroxide can form, such as with barium and strontium:
M (s) + O2 (g) $\,$$\,$ MO2 (s)

The peroxide ion (-O-O-) has two oxygen atoms, each with a weak covalent bond to the other, and a 1- charge (O22- overall). It does not form in this reaction with the smaller group 2 metals such as Be or Mg because they form highly polarizing ions.
These are small with a very high charge density which draws the 2- negative charge in the peroxide ion towards it with great strength (they polarize the ion). The peroxide ion becomes a stable O2- oxide ion and the weak O-O covalent bond breaks.
The heavier, less dense and less polarizing ions like Sr2+ and Ba2+ cannot do this, so the metal peroxide is stable.

• The reaction with oxygen to form a metal oxide occurs spontaneously in air and when the metals are heated by flame, showing a distinct colour.

• The following is a table of the flame colours observed in the group 1 and 2 elements:
 Group 1 elements Group 2 elements Lithium: red Beryllium: white/colourless Sodium: orange/yellow Magnesium: white/colourless Potassium: lilac Calcium: red/orange Rubidium: red Strontium: dark red Caesium: blue/violet Barium: green

• The reaction between the group 1 and 2 metals with chlorine can be written generally as:

• For a group 1 metal:
2M (s) + Cl2 (g) $\,$$\,$ 2MCl (s)

For a group 2 metal:
M (s) + Cl2 (g) $\,$$\,$ 2MCl (s)

• The solubility of many group 1 and 2 metal compounds have trends down the groups. Generally, group 1 metal compounds are more soluble than any group 2 analogue.
For example, you would predict that potassium hydroxide has greater solubility (in water at a fixed temperature) than calcium hydroxide.
• Solubility of the metal hydroxides increases down the group. This is true for both group 1 and group 2 hydroxides, and as said above the group 1 hydroxides are much more soluble
• Solubility of the group 2 metal sulfates decreases down the group. Barium sulfate is the product we look for in the common test for sulfate ions. The white precipitate is obvious as BaSO4 is extremely insoluble.
• Solubility of the group 2 metal carbonates generally decreases down the group. In group 1, solubility actually increases down the group.
Explaining the solubility trends is not necessary. In short, it’s because of two enthalpy measurements (lattice enthalpy and enthalpy of hydration) that change at different rates down the groups, depending on the different anions (sulfate, carbonate, etc) used... don’t worry about it!
• Introduction
Group 1 and group 2: a summary
a)
Intro to group 1 and 2.

b)
Properties of group 1 and group 2 metals.

c)
How do the properties change going down the group? (Trends in properties)

d)
Why do group 1 and 2 metals get more reactive down the group?

e)
Reactions of group 1 and 2 metals.

f)
Solubility of group 1 and 2 metals.

• 1.
Recall the trend in properties of the alkali metals.
Sort the alkali metals for the following properties:
i) From hardest to softest.
ii) From lowest to highest melting point.

• 2.
Recall and predict the reactions of the alkali metals.
In a school laboratory, a chunk of sodium metal is put in water.

i) Write a full word equation to show the reaction taking place.

ii) One of the products of the reaction can be identified by using universal indicator. What color will the solution become if universal indicator is added?

iii) After this, a student puts a chunk of potassium metal in a separate bath of water. Will this reaction be more or less vigorous? Explain why.