Group 1 and Group 7: Properties, Reactions and Reactivity Patterns

Chemical families show patterns

A group is a vertical column in the Periodic Table. Elements in the same group have similar chemical behaviour, so observations of a few members help us predict reactions of others. Two important families are the Group 1 alkali metals and the Group 7 halogens.

Their reactivity trends run in opposite directions: Group 1 becomes more reactive down the group, whereas Group 7 becomes less reactive down the group. Knowing the family and the position within it therefore helps us predict both the type of reaction and its likely vigour.

Recognising the Group 1 metals

The alkali metals are lithium, sodium, potassium, rubidium, caesium and francium. They occupy the far-left column of the Periodic Table, below hydrogen. Hydrogen is not an alkali metal.

Unlike many familiar metals, alkali metals are soft enough to cut with a knife. Their freshly cut surfaces are shiny and silvery, but quickly become dull in air. They have relatively low densities and low melting points compared with most metals. Lithium, sodium and potassium are less dense than water, so they float on its surface.

Moving down Group 1, the metals become softer and their melting points decrease. Density does not follow a perfectly regular increase: potassium is less dense than sodium. The reactivity trend, however, is clear—reactivity increases down the group.

Group 1 metals react with water to make an alkali

Lithium, sodium and potassium all react with water to produce hydrogen gas and a solution of the metal hydroxide:

metal + water → metal hydroxide + hydrogen

The hydroxide solution is alkaline, which explains the name alkali metals. Universal indicator turns towards blue or purple as the alkaline solution forms. Fizzing shows that gas is being released; in these reactions the gas is hydrogen.

The balanced equations are:

$$\mathrm{2Li(s) + 2H_2O(l) \rightarrow 2LiOH(aq) + H_2(g)}$$ $$\mathrm{2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)}$$ $$\mathrm{2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g)}$$

Here, $(s)$ means solid, $(l)$ liquid, $(aq)$ dissolved in water and $(g)$ gas. The reactions have the same pattern, but their vigour differs.

The reactions release heat. Sodium and potassium can become hot enough to melt during the reaction. Ignition is possible, but a flame is not needed to establish that a reaction has occurred.

These observations give the order lithium < sodium < potassium, from least to most reactive.

Moist air changes the metal surface

Group 1 metals react readily with both oxygen and water vapour in air. Reaction with oxygen produces a dull coating of oxygen-containing compounds, so the metal loses its shine: it tarnishes. Lithium, for example, forms lithium oxide:

$$\mathrm{4Li(s) + O_2(g) \rightarrow 2Li_2O(s)}$$

The moisture in air is water, so it can also react with the metal to form a metal hydroxide and hydrogen. A freshly cut surface therefore does not stay shiny when exposed to moist air.

Group 1 metals are stored under oil to keep oxygen and moisture away from their surfaces. This prevents unwanted reactions and helps preserve them before a demonstration.

Investigating the Group 1 reactivity trend safely

The reaction with water is a teacher demonstration, not a student experiment. The apparatus includes a water trough, universal indicator, small pieces of lithium, sodium and potassium, dry forceps and safety screens. Everyone wears eye protection, and students stay behind the screens at the teacher's viewing line.

The teacher removes a small metal piece from its oil using dry forceps, blots away the oil and places it on the water. Students record fizzing, movement, melting, any flame and the indicator colour change. The teacher confirms that the piece has completely disappeared and that the reaction has stopped before adding the next metal.

The teacher uses carefully limited pieces because larger pieces can react dangerously. Tools must be dry to prevent an unintended reaction with water. Students do not handle the metals or clear away scraps, and no other flames should be nearby because hydrogen is flammable.

Compare several observations together. Increasingly vigorous fizzing and movement support increasing reactivity from lithium to sodium to potassium. Time to disappearance can also be recorded, but it is not a fair numerical comparison when piece sizes differ. Melting changes the metal's contact with water too. Keep the water temperature and observation conditions comparable. This demonstration establishes a qualitative trend rather than an exact rate ratio.

Group 1 metals react with chlorine to form chlorides

When heated with chlorine gas, lithium, sodium and potassium react vigorously to form metal chlorides. These are white solid salts. The reaction becomes more vigorous down Group 1.

metal + chlorine → metal chloride

For the first three alkali metals:

$$\mathrm{2Li(s) + Cl_2(g) \rightarrow 2LiCl(s)}$$ $$\mathrm{2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)}$$ $$\mathrm{2K(s) + Cl_2(g) \rightarrow 2KCl(s)}$$

Notice that chlorine is written as $\mathrm{Cl_2}$, but each chloride formula contains one metal atom for each chlorine atom. The coefficients balance the numbers of atoms on both sides. Chlorine is toxic, so these reactions require specialist teacher-controlled arrangements, not the water-demonstration setup.

Recognising halogens in different states

Group 7 contains the halogens: fluorine, chlorine, bromine, iodine and astatine. They are reactive non-metals and are toxic. Their physical states and colours differ, even though their chemical behaviour follows a common pattern.

From fluorine to iodine, the colours generally become darker and melting and boiling points increase. This helps explain the change from gases to a liquid and then a solid at room temperature. Iodine's purple vapour is different from its grey-black solid and its brown aqueous solution.

Astatine lies below iodine. From the group pattern, we predict a dark solid at room temperature and lower reactivity than iodine. This is a prediction based on its position, rather than a classroom observation.

Halogens and halide salts are different substances

Halogens are diatomic: their molecules contain two atoms, as shown by the formulae $\mathrm{Cl_2}$, $\mathrm{Br_2}$ and .

A halogen atom can gain one electron to form a halide ion with a charge of $-1$. For example, chlorine forms chloride, $\mathrm{Cl^-}$. Potassium chloride contains chloride ions; it does not contain chlorine molecules. Similarly, potassium bromide contains bromide ions and potassium iodide contains iodide ions.

Halogens react with Group 1 metals to form metal halides. Chlorine produces chlorides, bromine produces bromides and iodine produces iodides. For example, sodium reacts with bromine to form sodium bromide:

$$\mathrm{2Na(s) + Br_2(l) \rightarrow 2NaBr(s)}$$

The corresponding reaction with iodine produces sodium iodide, $\mathrm{NaI}$. For the same metal under comparable conditions, these reactions generally become less vigorous down Group 7. This contrasts with the increasing reactivity down Group 1.

The difference between a halogen and a halide matters because displacement experiments mix a halogen solution with a solution containing halide ions.

A more reactive halogen displaces a less reactive one

Chlorine displaces bromine from potassium bromide solution:

chlorine + potassium bromide → potassium chloride + bromine

$$\mathrm{Cl_2(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_2(aq)}$$

The mixture becomes orange because bromine is released. Chlorine also displaces iodine from potassium iodide, giving a brown solution. Bromine displaces iodine from potassium iodide, but cannot displace chlorine from potassium chloride. Iodine displaces neither chlorine nor bromine from their salts.

These results establish chlorine > bromine > iodine in reactivity. A halogen does not displace itself, and a less reactive halogen cannot displace a more reactive one.

The rule also applies to other metal halides in solution. For example, chlorine can displace bromine from sodium bromide as well as potassium bromide: it is the bromide ions that take part in the change.

Investigating halogen displacement on a spotting tile

Use a spotting tile, labelled droppers, dilute chlorine water, bromine water and iodine solution, and potassium chloride, bromide and iodide solutions. Label three rows for the halogens and three columns for the halide solutions so that all nine combinations can be tested.

Add three drops of the appropriate halide solution to each well, then three drops of the appropriate halogen solution. Wait one minute and record the colour against a white background. Keep drop counts, observation time, lighting and halide concentrations the same, and use fresh halogen solutions.

Prepare control wells containing each halogen solution diluted with water, and each halide solution diluted with water. Use the same drop counts as in the reaction wells. The halide controls should be colourless. Dilute chlorine water may be very pale green or appear colourless; bromine water is pale orange and iodine solution pale brown. Compare mixtures with these controls to distinguish a reaction from simple dilution. Repeat an unclear result with fresh drops.

After the fixed observation time, record whether a new halogen colour has appeared. Chlorine producing orange bromine from bromide shows that chlorine is more reactive than bromine. Bromine producing brown iodine from iodide shows that bromine is more reactive than iodine. Together, these observations establish the order, rather than merely showing that some mixtures change colour.

Wear splash-resistant eye protection, avoid skin contact and never sniff the wells. Chlorine and bromine give hazardous vapours: preparation and dispensing are staff tasks in a fume cupboard. Use only the teacher-issued dilute solutions and keep bottles closed. After recording the colours, collect the mixtures and rinses in the labelled halogen-waste container.

Using position to make predictions

First identify the group, then apply its reactivity trend. Rubidium lies below potassium in Group 1, so we predict a more vigorous reaction with water, producing rubidium hydroxide and hydrogen. We also predict that rubidium reacts with chlorine to form rubidium chloride. These predictions are not reasons to attempt the reactions in class.

For Group 7 displacement, compare the two halogens involved. Bromine lies below chlorine, so bromine cannot displace chlorine from a chloride. It lies above iodine, so it can displace iodine from an iodide.

Position therefore helps predict the kind of product from the family pattern and relative reactivity from the trend within that family. It does not tell us an exact reaction time: observed speed also depends on conditions such as temperature, concentration and surface area.

Group 1

• Soft metals; shiny when freshly cut, but tarnish rapidly in moist air.
• Relatively low densities and melting points; Li, Na and K float on water.
• Melting points decrease down the group; reactivity increases: Li < Na < K.
• Oxygen causes surface oxidation; water vapour reacts too. Store under oil.
• Metal + water → metal hydroxide + hydrogen. Universal indicator turns blue/purple.
• Lithium fizzes without melting; sodium melts and moves; potassium reacts more vigorously and may give a lilac flame.
• Metal + chlorine → metal chloride, a white solid; reaction becomes more vigorous down Group 1.

$$\mathrm{2Na + 2H_2O \rightarrow 2NaOH + H_2}$$