# Group 2 and Group 7

The elements in group 2 all have alkaline hydroxides, which is why the common name for this group is the alkaline metals. The group 2 elements have the following properties:

• High melting and boiling points.
• A giant metallic structure with strong forces between positive and negative ions.
• They are light metals with low densities which form colourless compounds.
• They all have their outermost electron in the $s$ sub-shell, with two electrons more than the electron configuration of a noble gas. Due to this configuration, alkaline metals are strong reducing agents.

The reaction for a group 2 alkaline is:

$$M \rightarrow M^{2+} + 2e{-}$$

Reactivity increases down the group due to the increasing ease of losing electrons.

### Reaction with oxygen

The group 2 elements react vigorously with oxygen in a redox reaction, forming an oxide with the general formula $MO$ where $M$ is the group 2 element. An example reaction is shown below:

$$2Ca_{(s)} + O_{2(l)} \rightarrow 2CaO_{(s)}$$

In this reaction, the $Ca$ is oxidised from 0 to +2. The $O$ has been reduced from 0 to -2.

### Reaction with water

Group 2 elements react with water to form hydroxides with the general formula $M(OH)_{2}$, along with hydrogen gas. For example:
$$Ca_{(s)} + 2H_{2}O_{(l)} \rightarrow Ca(OH)_{2(aq)} + H_{2(g)}$$

This is also a redox reaction with $Ca$ being oxidised from 0 to +2 and one hydrogen being reduced from +1 to 0. The further down the group, the more vigorously they react.

### Reaction with oxides and hydroxides

Group 2 oxides and hydroxides, formed with the reaction with oxygen and water are bases. These can neutralise acids to form a salt and water. For example:

$$MgO_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_{2(s)} + H_{2}O_{(l)}$$

This is not a redox reaction however as the oxidation numbers remain unchanged.

### Reaction of group 2 oxides with water

Group 2 oxides react with water to form a solution of metal hydroxides. These hydroxides have a typical pH of 10-12. Such reaction is:

$$MgO_{(s)} + H_{2}O_{(l)} \rightarrow Mg(OH)_{2(aq)}$$

### Group 2 hydroxides

Group 2 hydroxides dissolve in water to form alkaline solutions. The solubility of the hydroxides increases down the group. The greater the solubility of the hydroxide, the more $OH^{-}$ ions are released into the solution and the more alkaline it becomes.

$$Ca(OH)_{2(aq)} + aq \rightarrow Ca^{2+}_{(aq)} + 2OH^{-}_{(aq)}$$

### Decomposition of group 2 carbonates

Group 2 carbonates can be decomposed by heat, forming an oxide and carbon dioxide. This type of reaction is called thermal decomposition. The further down the group, the more difficult it is to decompose.

$$MgCO_{3(s)} \rightarrow MgO_{(s)} + CO_{2(g)}$$

The further down the group, the more difficult it is to decompose the group 2 hydroxides.

### Uses of group 2 hydroxides

The alkaline nature of the group 2 hydroxides means that they are often used in order to combat acidity. Examples include:

• Calcium hydroxide, $Ca(OH)_{2}$ which is commonly used by farmers to neutralise acid soils.
• Magnesium hydroxide, $Mg(OH)_{2}$ is used as milk of magnesia to relieve indigestion by neutralising stomach acid.

## Group 7 Elements

The group 7 elements: fluorine ($F$), chlorine ($Cl$), bromine ($Br$), iodine ($I$) and astatine ($At$) are generally known as the halogens. They exist as simple diatomic molecules, with weak van der Waals forces. This gives the halogens low melting and boiling points, which increase down the group due to the increased number of electrons.

Halogens have one electron less than the noble gas configuration making them oxidising agents as they remove electrons from a reaction. The oxidising power is a measure of the strength with which a halogen is able to capture an electron, forming a halide ion.

Down the group halogens become less reactive due to:

• Increased electron shielding
• While the nuclear charge also increases, the effect is largely cancelled out by the shielding and the greater atomic radius.

### Halogen Redox Reactions

The reactivity of halogens can be determined through displacement reactions where halide ions are mixed with aqueous solutions of halogens. A more reactive halogen is able to oxidise and displace a halide of a less reactive halogen.

Bromine is able to oxidise $I^{-}$ as it is the more reactive halogen.

$$Br_{2(aq)} + 2I^{-}_{(aq)} \rightarrow 2Br^{-}_{(aq)} + I_{2(aq)}$$

This reaction is a redox reaction as the $Br$ has been reduced from 0 to -1 while the $I$ has been oxidised from -1 to 0. $Br_{2}$ is unable to oxidise $Cl^{-}$ ions as chlorine is a more reactive halogen.

The result of the reactions can be determined by looking at the colour change. This colour is determined by the halogen in the solution. Chlorohexane can be added to make this change more visible.

Halogen Water Cyclohexane
$Cl_{2}$ Pale green Pale green
$Br_{2}$ Orange Orange
$I_{2}$ Brown Violet

## Disproportionation of chlorine

A disproportionation reaction is where the same element is both oxidised and reduced.

#### In water

Small amounts of chlorine is added to drinking water in order to kill bacteria making it safer to drink. In water, the chlorine forms hydrochloric acid $HCl$, and chloric (I) acid $HClO$.

$$Cl_{2(aq)} + H_{2}O_{(l)} \rightarrow HClO_{(aq)} + HCl_{(aq)}$$

#### In aqueous sodium hydroxide

Household bleach is formed when aqueous sodium hydroxide reacts with dilute sodium hydroxide at room temperature.

$$Cl_{2(aq)} + 2NaOH_{(aq)} \rightarrow NaCl_{(aq)} + NaClO_{(aq)} + H_{2}O_{(l)}$$

## Use of halide compounds

Ionic halides have a halide ion, $X^{-}$ and form compounds with group 1 or group 2 elements. Such compounds are:

• Sodium chloride, $NaCl$ is used as table salt
• Sodium fluoride, $NaF$ is added to toothpaste to combat tooth decay.
• Calcium fluoride, $CaF_{2}$ is used to make lenses to focus infrared light.

### Detecting halide ions

Halide ions can be detected by adding silver nitrate, $AgNO_{3(aq)}$ to the reaction mixture. The process is as follows:

1. The unknown halide is dissolved into water.
2. An aqueous solution of silver nitrate, $AgNO_{3(aq)}$ is added.
3. The silver ions, $Ag^{+}_{(aq)}$ react with any halide ions $X^{-}_{(aq)}$ forming a silver halide precipitate, $AgX_{(s)}$ .
4. These precipitates are coloured and can be examined to find the halide present.

The equation for this reaction is:
$$AgNO_{3(aq)} + MX_{(aq)} \rightarrow AgX_{(s)} + MNO_{3(aq)}$$

The ionic equation is:
$$Ag^{+}_{(aq)} + X^{-}_{(aq)} \rightarrow AgX_{(s)}$$

It can be difficult to determine the colour of the solution therefore aqueous ammonia $NH_{3(aq)}$ can be added. Each halide has a different solubility in ammonia therefore this can be useful in determining the solubility.

Halide $AgNO_{3(aq)}$ Dilute $NH_{3}$ Concentrated $NH_{3}$
$Cl^{-}$ White precipitate Soluble Soluble
$Br^{-}$ Cream precipitate Insoluble Soluble
$I^{-}$ Yellow precipitate Insoluble Insoluble

This type of reaction is a precipitation reaction as a precipitate is formed.