Group 1

Group 1 in the periodic table is called the alkali metals. The reason they are called this is due to the fact that they in combination with water create a alkaline (basic) solution. The alkali metals of group 1 are more reactive than the alkaline earth metals of group 2. We will mention this further in the next article.

Group 1 contains the elements:

  • Hydrogen, H
  • Lithium, Li
  • Sodium, Na
  • Potassium, K
  • Rubidium, Rb
  • Caesium, Cs
  • Francium, Fr 


All members of group 1 have one valence electron. This makes them very reactive, since they can reach noble gas structure by donating an electron to elements in its vicinity, starting a number of different chemical reactions. Since they very easily donate electrons, we can call them very weakly electronegative, or even electropositive. The further down we go in the group, the more reactive the elements become. The reason for the increased reactivity is the increased atomic radius, which makes the outermost electron attach very weakly due to large distance to the atomic nuclei. Potassium therefore reacts more strongly (faster, more explosively, with more elements) than lithium, which is further up in the group.

All atoms except for hydrogen are metallic, which means they are very soft, has a metallic surface, and conducts heat and electricity. Since the alkali metals only has one electron, the metal bond is relatively weak, which makes them soft compared to other metals.

Hydrogen is the exception

Hydrogen is part of group 1, but is not an alkali metal. Hydrogen as an element has one electron, but only needs one extra electron to fill its outer electron shell. This means that the hydrogen atom has properties which is similar to both group 1 (one valence electron), and the halogens of group 17 (which only need one more electron to reach a noble gas structure). In its common state, hydrogen forms a diatomic gas, and thus doesn't look like the alkali metals to any larger extent.


The alkali metals react with pretty much any substance that can accept electrons. Examples of this is water and halogens (elements from group 17).

When sodium reacts with water, the following reaction takes place:

\(\mathrm{2Na(s) + 2H_2O(l) \longrightarrow \:2NaOH(aq) + H_2(g)}\)

Due to the sodium hydroxide created in the reaction, the resulting solution is basic (alkaline) after the reaction.

One should also mention that the further down in the group we go, the more explosive the reaction with water becomes. The reaction between sodium and water is so exothermic (giving energy to the surroundings) that the hydrogen gas created can catch fire. This makes the sodium metal appear to burn, or creates explosions which throws metal fragments around. The reactions become worse further down the group. The reaction between potassium and water always results in the hydrogen gas to catch fire. Here is a video that shown how it looks when alkali metals react with water.

In reaction with halogens, salts are created.

\(\mathrm{2Na(s) + Cl_2(g) \longrightarrow \:2NaCl(s)}\)

Here is a video of the reaction between sodium and chlorine.

Abundance in nature

Since all alkali metals react with pretty much anything in their surroundings, they are not present in pure form in nature. They instead exist as dissolved ions in water (for example NaCl in ocean water) or as solid minerals in mountains and in the ground. Sodium and potassium ions are extremely important for us, as the movement of these ions across the cell membrane creates the signals of our nervous system.

The next article in the series is about group 2.

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