Group 1

Who is in Group 1?

The first column of the periodic table contains the alkali metals: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). (Hydrogen is sometimes placed in Group 1 in the periodic table but is NOT an alkali metal — it behaves quite differently.)

All Group 1 atoms have one outer-shell (valence) electron:

• Li: 2,1
• Na: 2,8,1
• K: 2,8,8,1
• Rb: 2,8,18,8,1
• Cs: 2,8,18,18,8,1
• Fr: 2,8,18,32,18,8,1

This single outer electron is at the heart of all Group 1 chemistry — they react by LOSING it to form a +1 ion (e.g. Na → Na⁺ + e⁻).

Physical properties of Group 1 metals.

1. Soft. All alkali metals are soft enough to cut with a knife — Na and K cut like cheese; Cs is almost like butter. Compare with hard metals like iron and copper.

2. Low density. Lithium (0.53 g/cm³), sodium (0.97), potassium (0.86) all FLOAT on water (density 1.00 g/cm³). Caesium is denser (1.93) but still much less than iron (7.87) or copper (8.96). Reason: the atoms are large and the metallic bonding is weak.

3. Low melting and boiling points. Mp/bp values (°C):

   Element	Mp	Bp
   Li	181	1342
   Na	98	883
   K	64	759
   Rb	39	688
   Cs	28	671

   Compare with iron (mp 1538 °C). Group 1 metals have WEAK metallic bonding — only one delocalised electron per atom contributes to the 'sea of electrons'.

4. Shiny when freshly cut. A newly cut surface looks bright silver but TARNISHES rapidly (within seconds) as it reacts with oxygen and water vapour in the air to form a dull oxide/hydroxide coating.

5. Good conductors of heat and electricity. Like all metals, they have a sea of mobile delocalised electrons that carry charge and energy.

Storage. Because they tarnish so quickly and react violently with water, alkali metals are stored UNDER OIL (paraffin oil). The oil keeps both air and water from the metal surface. Pieces are handled with tweezers and cut with a knife on a dry tile.

Group 1 metals react with water to produce an alkaline metal hydroxide and hydrogen gas. The general equation:

$2\text{M(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{MOH(aq)} + \text{H}_2\text{(g)}$

The reaction is highly exothermic — heat released raises the local temperature; can ignite the hydrogen.

Lithium.

$2\text{Li(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{LiOH(aq)} + \text{H}_2\text{(g)}$

Observations: Li floats, fizzes steadily (bubbles of H₂), gradually dissolves over a few minutes. NO ignition — the reaction is too gentle to heat the H₂ enough.

Sodium.

$2\text{Na(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{NaOH(aq)} + \text{H}_2\text{(g)}$

Observations: Na floats, fizzes vigorously, MELTS into a silvery ball (low mp 98 °C; exothermic reaction provides enough heat). The molten ball rolls rapidly across the surface, propelled by the H₂ gas. Sometimes the H₂ ignites with a yellow-orange flame, ending in a small 'pop'.

Potassium.

$2\text{K(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{KOH(aq)} + \text{H}_2\text{(g)}$

Observations: K floats, fizzes very violently, melts immediately and IGNITES — burning with a LILAC flame (the characteristic colour of K⁺ ions). The piece may rush across the surface and end in a small explosion. Reaction over in 10-20 seconds.

Rubidium and caesium react EXPLOSIVELY with water — dangerous to demonstrate without a protective enclosure. Both ignite immediately on contact.

Detecting the alkali formed. The MOH product is a STRONG ALKALI (pH 13-14 in solution):

• Turns red litmus paper BLUE.
• Turns universal indicator PURPLE/dark blue.
• Reacts with acids to give a salt + water (neutralisation).

If you drop sodium into water that already contains a few drops of universal indicator, the indicator turns purple in a 'cloud' around the metal as NaOH dissolves.

Detecting the hydrogen. Hold a lighted splint near the H₂ being evolved — it produces a 'SQUEAKY POP' (a small explosion as H₂ burns with O₂ in air).

Reactivity in Group 1 is judged by how readily the atom loses its single outer-shell electron to form M⁺. Going down the group from Li to Cs, three structural changes occur — and together they make the electron easier to lose.

1. Atomic radius INCREASES. Each successive element has an extra electron shell:

• Li: 2 shells (outer in shell 2)
• Na: 3 shells (outer in shell 3)
• K: 4 shells (outer in shell 4)
• Rb: 5 shells (outer in shell 5)
• Cs: 6 shells (outer in shell 6)

The outer electron is therefore FURTHER from the nucleus. Atomic radii: Li 152 pm → Cs 265 pm — almost double.

2. Shielding INCREASES. The inner filled shells partially shield the outer electron from the positive nuclear charge — they act as a 'screen'. Going down the group, MORE inner shells means MORE shielding:

• Li: 1 inner shell shields the outer electron.
• Na: 2 inner shells shield.
• K: 3 inner shells.
• Cs: 5 inner shells.

3. Nuclear charge INCREASES (but is outweighed). Each element has more protons (Li 3p → Cs 55p). You might think this would make the outer electron HARDER to remove. But the increases in atomic radius and shielding more than offset this — the EFFECTIVE attraction on the outer electron actually DECREASES.

Net effect. Going down the group, the attraction between the nucleus and the outer electron WEAKENS. So the outer electron is HELD LESS TIGHTLY and LOST MORE EASILY → reactivity INCREASES.

This is shown clearly by the FIRST IONISATION ENERGIES (energy needed to remove the outer electron):

Lower IE → easier to lose outer electron → more reactive.

Linking to observed reactivity with water. As we saw, Li reacts gently; Na vigorously; K ignites. The trend in violence MATCHES the trend in first ionisation energy. The most reactive Group 1 metals (Cs, Fr) are extremely vigorous, sometimes explosive.

The opposite trend in Group 7. Note the contrast: Group 7 reactivity DECREASES down the group. Why? Because Group 7 reacts by GAINING an electron (not losing one) → the trend in 'ease of gaining an electron' goes the OTHER way down the group. Same atomic-structure reasoning, opposite reactive process.

Beyond the reaction with water, Group 1 metals show predictable behaviour with other substances.

Reaction with oxygen. Group 1 metals burn in air (or oxygen) with characteristic flame colours, forming a metal OXIDE:

$4\text{Li(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{Li}_2\text{O(s)} \text{ (red flame)}$ $4\text{Na(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{Na}_2\text{O(s)} \text{ (orange-yellow flame)}$ $4\text{K(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{K}_2\text{O(s)} \text{ (lilac flame)}$

The oxides are WHITE solids and react with water to form the same alkali as before:

$\text{Na}_2\text{O(s)} + \text{H}_2\text{O(l)} \rightarrow 2\text{NaOH(aq)}$

Note: In practice, Na also forms Na₂O₂ (peroxide) and K forms KO₂ (superoxide) under excess O₂, but Edexcel 4CH1 only requires the simple oxide M₂O.

Reaction with halogens. Group 1 metals react vigorously with Group 7 elements (halogens) to form IONIC halides:

$2\text{Na(s)} + \text{Cl}_2\text{(g)} \rightarrow 2\text{NaCl(s)}$ $2\text{K(s)} + \text{Br}_2\text{(l)} \rightarrow 2\text{KBr(s)}$

The reactions are highly exothermic. The metal loses 1 electron; the halogen gains 1 electron. The resulting compound is a typical IONIC SALT — white crystalline solid, soluble in water, high mp, conducts when molten or dissolved.

Flame test colours. The characteristic colours imparted by Group 1 ions in a flame are used for identification:

Method for flame test. Use a clean nichrome (or platinum) wire — dip in concentrated HCl, then into the solid sample. Hold in the HOTTEST part of a BLUE Bunsen flame (air-hole fully open). The HCl cleans the wire and volatilises the metal ion; the heat excites electrons in the ion, which emit characteristic colours as they return to the ground state. Always clean the wire between samples to avoid contamination — even traces of Na drown out K's lilac with strong yellow (cobalt-blue glass can filter out the yellow to confirm K).

Francium (Fr) is the heaviest Group 1 metal. It is RADIOACTIVE — the most stable isotope (²²³Fr) has a half-life of only ~22 minutes. The total amount of Fr on Earth at any given time is estimated at less than 30 g, so bulk samples have never been studied. We predict its properties by EXTRAPOLATING from the observed trends in Li → Cs.

Predicted property 1 — Extreme reactivity with water.

Continuing the trend Li < Na < K < Rb < Cs, francium would react EVEN MORE VIOLENTLY than caesium — almost certainly explosively. Equation predicted:

$2\text{Fr(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{FrOH(aq)} + \text{H}_2\text{(g)}$

Reason: Fr's outer electron sits in the 7th shell, furthest from the nucleus, with the most inner-shell shielding of any Group 1 element. Lowest 1st ionisation energy → outermost electron most easily lost.

Predicted property 2 — Very low melting point.

Mp trend: Li 181 → Cs 28 °C. Extrapolation suggests Fr might melt at ~20-25 °C — potentially a LIQUID at warm room temperature. Reason: weak metallic bonding (only one outer electron contributes to the electron sea), and the largest atomic radius gives the weakest bonds.

Predicted property 3 — Soft, low density.

Like Cs but even softer. Density predicted ~2.0-2.5 g/cm³ — still low for a metal, continuing the trend Cs 1.93.

Predicted property 4 — Forms strongly alkaline FrOH.

By extension, Fr would form a STRONG alkali FrOH in water — more alkaline than KOH or CsOH. Reacts with acids to form a salt + water.

Predicted property 5 — Brilliant flame test colour.

Extrapolating from the colour trend, francium's flame test would be expected to be a redder/more orange colour relative to Cs's blue — though it has never been measured in bulk quantity.

Limitation — radioactivity. Francium isotopes decay by alpha or beta emission. Bulk samples are NOT made. Predictions are based on TRENDS — never verified by experiment on macroscopic amounts.

Strategy for exam questions about Fr. When asked to predict properties of Fr, always:

1. State the relevant Group 1 trend (reactivity, mp, density, etc.).
2. Apply the trend to Fr (more reactive, lower mp, etc.).
3. Justify briefly with atomic structure (more shells, more shielding, outer e⁻ further out).

Examiners reward atomic-structure REASONING, not just 'because it's at the bottom'.