Group 1: The Alkali Metals

The metals in Group 1 — lithium, sodium, potassium and those below them — are the alkali metals, and they behave nothing like the everyday metals most people picture. They are soft enough to cut with a knife, they are so light that the first three float on water, and above all they are astonishingly reactive, bursting into life the instant they meet water or air. In this lesson, part of Topic C4 of OCR Gateway Combined Science A, we focus on their reactions — with water and with oxygen — and on the striking way their reactivity increases as you go down the group, which can be explained by how easily each atom loses its outer electron. (You met the structure of the periodic table and how it is arranged in Topic C2; here the emphasis is on reactions and prediction rather than on re-teaching periodicity.)

By the end of this lesson you should be able to describe the properties of the alkali metals, write word and balanced symbol equations for their reactions with water and oxygen, describe and explain the reactivity trend down Group 1 in terms of electron loss, and use that trend to predict how an unfamiliar Group 1 metal will behave.

This lesson builds AO1 recall of the properties and reactions of the alkali metals, AO2 application when you write balanced symbol equations for their reactions with water and oxygen, and AO3 analysis when you explain the reactivity trend by electron loss and use it to predict an unfamiliar metal's behaviour.

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Properties of the Alkali Metals

The Group 1 metals share a distinctive set of physical and chemical properties that set them apart from ordinary metals:

• they are soft — sodium can be sliced with a knife — and have low densities, so lithium, sodium and potassium all float on water;
• they have relatively low melting points for metals, which fall further down the group;
• they are very reactive, so they must be stored under oil to keep air and water away from them;
• they always form +1 ions by losing their single outer electron, and their compounds are white ionic solids that dissolve to give colourless solutions.

That last point is the key to their chemistry. Each alkali-metal atom has exactly one electron in its outer shell, and reacting almost always means losing that one electron to form a stable $+1$+1 ion. Because every member of the group reacts in the same way — by losing that single outer electron — they all show similar chemical properties, which is exactly why they are grouped together in the first place. The differences between them are differences of degree (how vigorously they react), not of kind, and those differences follow the clear trend we come to below.

Their extreme reactivity also explains how they must be handled and stored. Left exposed to air, an alkali metal would react with both oxygen and water vapour, so a freshly cut piece is dull within seconds and is kept submerged in oil to seal it off. In the laboratory only small pieces are used, behind a safety screen, precisely because the reaction with water is so energetic. This practical caution is itself evidence of where these metals sit — right at the top of the reactivity series.

Exam Tip: The name "alkali metal" is a clue worth remembering: these metals react with water to make alkalis (metal hydroxide solutions). A common misconception is that all metals are hard and dense like iron — the alkali metals are the opposite, being soft and light enough to float.

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Reaction with Water

The reaction that best shows off the alkali metals is their reaction with water. They react vigorously with cold water to produce a metal hydroxide — an alkali — and hydrogen gas:

$\text{metal} + \text{water} \rightarrow \text{metal hydroxide} + \text{hydrogen}$metal+water→metal hydroxide+hydrogen

For sodium, the word and balanced symbol equations are:

$\text{sodium} + \text{water} \rightarrow \text{sodium hydroxide} + \text{hydrogen}$sodium+water→sodium hydroxide+hydrogen $2\text{Na} + 2\text{H}_2\text{O} \rightarrow 2\text{NaOH} + \text{H}_2$2Na+2H2​O→2NaOH+H2​

The metal floats, fizzes and skims about on the surface as hydrogen is released. What differs from metal to metal is how vigorous this is, and that difference reveals the trend:

• Lithium reacts steadily, fizzing gently as it floats.
• Sodium reacts faster, melting into a shiny ball that darts across the surface.
• Potassium reacts so vigorously that the hydrogen produced ignites, burning with a lilac flame, and the reaction may spit.

The reaction with water always leaves an alkaline solution, because a soluble metal hydroxide is formed; a few drops of universal indicator would turn purple. Reading down the list, the reactions get more and more energetic, which is the headline pattern of the group: reactivity increases down Group 1.

Alkali metal	Reaction with cold water	Trend
Lithium	Steady fizzing, floats and moves slowly	least reactive
Sodium	Vigorous, melts into a ball, darts about	more reactive
Potassium	Very vigorous, hydrogen ignites (lilac flame)	most reactive (of these three)

Exam Tip: For any alkali metal + water, the products are metal hydroxide + hydrogen — never an oxide (that is the steam reaction of less reactive metals). Writing "sodium oxide" for sodium in cold water is a classic error.

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Reaction with Oxygen

The alkali metals also react readily with oxygen in the air to form metal oxides, which is why a freshly cut surface quickly loses its shine and tarnishes — and why the metals are kept under oil in the first place. The general pattern is:

$\text{metal} + \text{oxygen} \rightarrow \text{metal oxide}$metal+oxygen→metal oxide

For lithium, this gives lithium oxide:

$4\text{Li} + \text{O}_2 \rightarrow 2\text{Li}_2\text{O}$4Li+O2​→2Li2​O

and for sodium, sodium oxide:

$4\text{Na} + \text{O}_2 \rightarrow 2\text{Na}_2\text{O}$4Na+O2​→2Na2​O

Because each metal forms a $+1$+1 ion and oxygen forms an $\text{O}^{2-}$O2− ion, the oxide always has the formula $\text{M}_2\text{O}$M2​O — two metal ions for every oxide ion. The speed of tarnishing follows the same trend as the water reaction: potassium tarnishes fastest, lithium slowest, once again showing that reactivity increases down the group.

Exam Tip: The formula of a Group 1 oxide is $\text{M}_2\text{O}$M2​O (for example $\text{Na}_2\text{O}$Na2​O), because the $+1$+1 metal ion must balance the $2-$2− charge on the oxide ion. Balancing the equation then needs a 4 in front of the metal and a 2 in front of the oxide.

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Explaining the Trend: Reactivity Increases Down the Group

Why should potassium be more reactive than lithium when both have just one outer electron? The answer lies in how tightly that outer electron is held. Going down the group, each element has one more electron shell than the one above it, so:

• the single outer electron is further from the nucleus; and
• it is shielded from the positive nucleus by more inner shells of electrons.

Both effects weaken the pull between the nucleus and the outer electron. Since reacting means losing that outer electron, and a weakly held electron is lost more easily, the metals become more reactive as you go down the group. This is a lovely example of a reactivity trend being explained entirely by electronic structure.

flowchart TD
  A["Go DOWN Group 1"] --> B["One more electron shell"]
  B --> C["Outer electron further from nucleus<br/>and more shielded"]
  C --> D["Weaker attraction to the nucleus"]
  D --> E["Outer electron lost more easily"]
  E --> F["More reactive down the group"]

Exam Tip: To explain the Group 1 trend, always join up the full chain: more shells → outer electron further out and more shielded → weaker attraction → electron lost more easily → more reactive. Stating only "there are more shells" without linking it to losing the electron will not gain full marks.

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Worked Examples

Worked Example 1 — Predicting a Group 1 reaction