Guides · 03
Alkali Metals and Alkaline Earth Metals: A Guide to Groups 1 and 2
Groups 1 and 2 explained: properties of alkali and alkaline earth metals, reactivity trends, flame test colors, everyday uses and where hydrogen belongs.
The two columns on the far left of the periodic table hold some of the most reactive metals there are. Group 1, lithium, sodium, potassium, rubidium, caesium and francium, are the alkali metals. Group 2, beryllium, magnesium, calcium, strontium, barium and radium, are the alkaline earth metals. Both belong to the s-block, and the simple fact that their atoms carry only one or two electrons in the outermost shell explains most of how they behave.
Where the names come from
"Alkali" comes from an Arabic word for plant ashes. People once leached the ashes of burned wood or seaweed with water to make soap and glass, and those ashes were rich in strongly basic potassium and sodium carbonates. The metals later extracted from such substances became the alkali metals.
"Earth" was an old chemical term for oxides that survive strong heating without change. Lime (calcium oxide) and magnesia (magnesium oxide) were earths that dissolved in water to give basic solutions, hence alkaline earths. The metals themselves were first obtained in 1808, when Humphry Davy in England used electrolysis to isolate calcium, magnesium, strontium and barium, a year after he had done the same for sodium and potassium.
Group 1: shared properties
Every alkali metal atom has a single electron in its outer s orbital. That electron is far from the nucleus and screened by the inner electrons, so it is lost easily. As a result these elements almost always appear as +1 ions.
- They are soft. Sodium and potassium can be cut with a knife like cold butter. A fresh surface is silvery but tarnishes within seconds in air.
- They are light. Lithium has a density of about 0.53 g/cm³, the lowest of any metal, and floats on water. Sodium and potassium are also less dense than water.
- They melt easily. Lithium melts at roughly 180 °C, sodium at about 98 °C, potassium at about 63 °C and caesium at about 28 °C. Caesium would melt on a warm hand, but it must never be touched, because it reacts violently with moisture on skin.
- They are stored under oil or inert gas, because they react with oxygen and water vapor in the air.
Reactivity increases down the group
Alkali metals react with water to produce hydrogen gas and a hydroxide. For sodium:
2Na + 2H₂O → 2NaOH + H₂
The same reaction becomes far more violent further down the group. Lithium fizzes gently and slowly disappears. Sodium melts into a bead that races across the surface. Potassium ignites the hydrogen it releases, burning with a lilac flame. Rubidium and caesium react explosively.
The trend follows ionization energy. Each step down adds an electron shell, so the outer electron sits farther from the nucleus and is easier to remove. First ionization energies fall from about 5.39 eV for lithium to 5.14 eV for sodium, 4.34 eV for potassium and 3.89 eV for caesium.
Reactions with oxygen show another difference. Lithium mainly forms the oxide Li₂O, sodium tends to form the peroxide Na₂O₂, and potassium and the heavier metals form superoxides such as KO₂. Larger cations form more stable lattices with larger anions.
Francium sits at the bottom of group 1, but its longest-lived isotope has a half-life of only about 22 minutes. A visible amount has never been collected, so most of its properties are estimates based on its neighbors.
Group 2: shared properties
Alkaline earth metals have two electrons in the outer s orbital and form +2 ions. Removing two electrons costs more energy, so they are less reactive than the group 1 metal in the same period, and they are harder, denser and have higher melting points.
The trend within the group mirrors group 1. Beryllium barely reacts with water, magnesium reacts with hot water or steam, and calcium, strontium and barium react with cold water to release hydrogen.
Ca + 2H₂O → Ca(OH)₂ + H₂
Beryllium is the odd one out. Its atom is so small that it tends to form covalent rather than ionic bonds, and its oxide is amphoteric, reacting with both acids and bases. Beryllium compounds can also cause serious lung disease if inhaled as dust, so they are handled with care.
Flame test colors
When compounds of the alkali metals and some alkaline earth metals are held in a flame, each element gives a characteristic color. Heat excites electrons to higher energy levels, and as they fall back they emit light at specific wavelengths. Fireworks rely on the same effect.
| Element | Flame color |
|---|---|
| Lithium | crimson red |
| Sodium | yellow |
| Potassium | lilac |
| Rubidium | red-violet |
| Caesium | blue-violet |
| Calcium | orange-red |
| Strontium | bright red |
| Barium | pale green |
Sodium's yellow is so intense that it can hide other colors. Chemists look at potassium's lilac flame through cobalt-blue glass to filter out any sodium yellow. Beryllium and magnesium give no distinctive flame color; the dazzling white light of burning magnesium ribbon comes from the combustion itself, not from a flame test emission.
Everyday and industrial uses
- Lithium: rechargeable lithium-ion batteries in phones and electric cars, lightweight aluminium-lithium alloys, and lithium carbonate as a mood-stabilizing medicine.
- Sodium: table salt (sodium chloride), baking soda (sodium hydrogen carbonate), yellow sodium-vapor street lamps, and coolant in some nuclear reactors.
- Potassium: one of the three main plant nutrients and a key ingredient in fertilizer. In the body it is essential for nerve signals and heart rhythm.
- Rubidium and caesium: atomic clocks. The SI second is defined by a specific transition frequency of the caesium-133 atom.
- Magnesium: light alloys for car parts and laptop cases, and the central atom of chlorophyll.
- Calcium: bones and teeth, limestone, cement and plaster. "Hard" tap water is water rich in calcium and magnesium ions.
- Strontium: the red in fireworks and signal flares.
- Barium: barium sulfate is so insoluble that it is safe to swallow as an X-ray contrast agent, even though soluble barium compounds are toxic.
Where does hydrogen belong?
Hydrogen has a single electron, so most tables put it at the top of group 1. Yet hydrogen is a nonmetallic gas at room temperature and behaves very differently from the alkali metals.
The case for group 1 rests on its configuration, 1s¹, and its ability to form H⁺. On the other hand, like a halogen it can gain one electron to form the hydride ion H⁻ and it exists as a diatomic molecule, H₂, so some argue it should sit above group 17. Others point out that its outer shell is exactly half full, like carbon's, and suggest group 14.
That is why some tables place hydrogen above group 1 but color it as a nonmetal, while others float it by itself above the middle of the table. Rather than having one correct answer, the debate shows that hydrogen is unique and matches no group perfectly. Theorists have also predicted that under the extreme pressures deep inside planets like Jupiter, hydrogen may exist in a metallic state.
Diagonal relationships
A period 2 element often resembles the period 3 element diagonally below and to its right. This is called a diagonal relationship. Lithium shares several traits with magnesium rather than sodium: it forms a nitride, and its carbonate decomposes readily on heating. Beryllium resembles aluminium, most clearly in its amphoteric oxide. The usual explanation is that the ratio of charge to size becomes similar along the diagonal. Knowing these exceptions alongside the main group trends gives a much deeper picture of groups 1 and 2.