Trends in Chemical Properties of Group 1 Elements and Their Compounds

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Senior 4 · Chemistry

Trends in Chemical Properties of Group 1 Elements and Their Compounds

4 lessons

Lesson 1 of 4

The Alkali Metals: Occurrence and Physical Properties

25 min

Metals you can cut with a knife

Most metals you know, iron in a hoe, copper in wire, aluminium in a saucepan, are hard and strong. The metals of Group 1 are different. Sodium can be cut with a table knife, potassium is softer still, and caesium melts on a warm day. They are so reactive that they are never found free in nature and must be stored under oil or in sealed glass. Yet their compounds are everywhere: common salt, the potassium in plant fertilisers, the lithium in phone batteries. This unit uses the ideas of Units 2 to 5 to explain their chemistry. They are called the alkali metals because they react with water to form strongly alkaline solutions.

A lump of sodium metal sealed in a small glass vial to keep it away from moist air. Sodium is soft enough to cut with a knife; a freshly cut surface is shiny but quickly dulls. Photo: Hi-Res Images of Chemical Elements, CC BY 3.0, via Wikimedia Commons

Definition

Alkali metals: the elements of Group 1, lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr). Each has one electron in its outer s orbital, $n s^{1}$ , and forms a 1+ ion, $\mathrm{M^{+}}$ , with the configuration of the noble gas before it. Hydrogen is placed above Group 1 but is a non-metal and is not an alkali metal.

Where they occur

Because they react so readily, the alkali metals occur only as compounds, mostly as $\mathrm{M^{+}}$ ions in salts. Sodium and potassium are among the ten most abundant elements in the Earth's crust. Sodium is found as rock salt (halite, NaCl) and dissolved in sea water and salt lakes; potassium in minerals such as sylvite (KCl) and carnallite, and in feldspar and clay minerals in soils. Lithium comes from minerals such as spodumene and from salty brines. Rubidium and caesium are rare, and francium is radioactive: only tiny amounts exist at any moment.

Sodium is extracted by the electrolysis of molten sodium chloride: sodium forms at the cathode and chlorine at the anode. No chemical reducing agent is strong enough to take electrons back from $\mathrm{Na^{+}}$ cheaply.

Physical properties down the group

Table 1.1: The alkali metals

Property	Li	Na	K	Rb	Cs
Atomic number	3	11	19	37	55
Outer configuration	$2 s^{1}$	$3 s^{1}$	$4 s^{1}$	$5 s^{1}$	$6 s^{1}$
Metallic radius / pm	152	186	227	248	265
Ionic radius of $\mathrm{M^{+}}$ / pm	76	102	138	152	167
First ionisation energy / kJ mol⁻¹	520	496	419	403	376
Electronegativity	1.0	0.9	0.8	0.8	0.8
Melting point / °C	181	98	63	39	28
Boiling point / °C	1342	883	759	688	671
Density / g cm⁻³	0.53	0.97	0.86	1.53	1.93

Radius increases down the group: one more shell each time. Each $\mathrm{M^{+}}$ ion is much smaller than its atom because it has lost its whole outer shell.
First ionisation energy decreases: the outer s electron is further from the nucleus and more shielded, so it is easier to remove. This is the key to the group's chemistry.
Melting and boiling points decrease: each atom gives one electron to the metallic bond, but the ions get larger, so the attraction between the ions and the electron sea weakens (Unit 3).
Density generally increases, because mass rises faster than volume. Potassium is the exception: it is less dense than sodium. Lithium, sodium and potassium are less dense than water and float on it.
All are soft, silvery and good conductors of heat and electricity, with only one delocalised electron per atom.

Pearls of potassium metal under paraffin oil; the larger one is about 0.5 cm across. The oil keeps out air and moisture, which would attack the metal at once. Photo: Hi-Res Images of Chemical Elements, CC BY 3.0, via Wikimedia Commons

Worked example 1: predicting francium

Use Table 1.1 to predict, with reasons, three physical properties of francium (Z = 87), the element below caesium.

Solution. Follow the trends one step further and explain each.

Melting point below 28 °C, perhaps a little above room temperature: the $\mathrm{Fr^{+}}$ ion is larger still, so the metallic bonding is weaker.
First ionisation energy below 376 kJ mol⁻¹: the 7s electron is further out and more shielded. (The measured value is about 393 kJ mol⁻¹, slightly higher than caesium's, an effect of the very large nuclear charge that this simple model does not include.)
Atomic radius larger than 265 pm, with one more shell.

Francium's properties are hard to check because it is so radioactive that no visible sample has ever been made.

Worked example 2: using electronic structure

Explain why the alkali metals form only 1+ ions in their compounds, using the ionisation energies of sodium: first 496, second 4562 kJ mol⁻¹.

Solution. Sodium is $1s^2\,2s^2\,2p^6\,3s^1$ . The first electron comes from 3s, far from the nucleus and shielded by 10 inner electrons, so it is removed easily. The second would come from the full 2p subshell, much closer to the nucleus and less shielded: it costs about nine times as much energy. The extra energy released by forming compounds of $\mathrm{Na^{2+}}$ could never repay this, so sodium is always +1 in its compounds.

Warning

Common mistakes. (1) Saying the melting point falls because "the atoms are heavier": it falls because the ions are larger and the metallic bond weaker. (2) Saying density increases steadily: potassium is less dense than sodium. (3) Calling hydrogen an alkali metal.

Safety. Alkali metals are stored under oil and handled only with tongs or tweezers, never fingers: the moisture on skin reacts with them and makes a caustic alkali. Cut them on a dry tile, use only pieces the size of a grain of rice, and return unused metal to the oil at once. Never put waste sodium or potassium in a sink or bin.

Tip

Group 1 trends come from one fact: one outer s electron, further from the nucleus each step down. Larger atoms → lower ionisation energy → more reactive; larger ions → weaker metallic bonding → lower melting point.

Cognito: Group 1 Elements - Alkali Metals (6 min). Watch for the properties shared by the whole group and the explanation of why reactivity increases down it, using the distance of the outer electron from the nucleus.

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Practice

Write the electron configuration of potassium and of the potassium ion.
Explain why the alkali metals are never found free in nature.
Describe and explain the trend in atomic radius down Group 1.
Explain why the melting point of lithium (181 °C) is much higher than that of caesium (28 °C).
Which Group 1 metal is the exception to the trend in density? Which metals float on water?
Explain why sodium is extracted by electrolysis rather than by heating its oxide with carbon.
Predict, with a reason, whether rubidium's first ionisation energy is greater or less than potassium's. Check your answer with Table 1.1.

Show answersHide answers

K: $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^1$ ; $\mathrm{K^{+}}$ : $1s^2\,2s^2\,2p^6\,3s^2\,3p^6$ .
They lose their single outer electron very easily and react quickly with oxygen, water and other substances, so they exist only as $\mathrm{M^{+}}$ ions in compounds.
It increases: each element has one more shell of electrons, further from the nucleus, with more inner shells shielding the outer electron.
Both have one delocalised electron per atom, but $\mathrm{Li^{+}}$ (76 pm) is much smaller than $\mathrm{Cs^{+}}$ (167 pm), so the electron sea is held much more strongly in lithium.
Potassium (0.86 g cm⁻³) is less dense than sodium (0.97). Lithium, sodium and potassium float on water.
Sodium is so reactive that its ions are very hard to reduce: carbon is not a strong enough reducing agent to turn $\mathrm{Na^{+}}$ back into the metal at practical temperatures, so electrolysis supplies the electrons directly at the cathode.
Less: rubidium's outer electron is in shell 5, further out and more shielded. Table 1.1: Rb 403, K 419 kJ mol⁻¹.

Quick check

Why do the melting points of the alkali metals decrease from lithium to caesium?

Key points

Group 1 (Li to Cs) all have an $n s^{1}$ outer electron and form $\mathrm{M^{+}}$ ions; they occur only in compounds and are stored under oil.
Down the group: atomic and ionic radius increase; first ionisation energy, electronegativity, melting and boiling points decrease; density generally increases (K is an exception).
Sodium is extracted by electrolysis of molten NaCl.
Handle with tongs, tiny pieces, goggles; never touch with bare hands.

Sources and credits

Image: Sodium (11 Na).jpg by Hi-Res Images of Chemical Elements, via Wikimedia Commons, licensed under CC BY 3.0.
Image: 0.5 grams lithium under argon.jpg by Hi-Res Images of Chemical Elements, via Wikimedia Commons, licensed under CC BY 3.0.
Image: Potassium-2.jpg from images-of-elements.com, via Wikimedia Commons, licensed under CC BY 3.0.
Radii, ionisation energies, melting and boiling points and densities are standard reference values, as listed in the CRC Handbook of Chemistry and Physics.
Video: GCSE Chemistry - Group 1 Elements - Alkali Metals, Cognito, embedded from YouTube.
The diagram, table, worked examples and Practice items were written and drawn by Ganzaa.

1Metals you can cut with a knife

2Where they occur

3Physical properties down the group

Metals you can cut with a knife

Where they occur

Physical properties down the group