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Reactions of Group 1 Elements with Water

All of Group 1 elements—lithium, sodium, potassium, rubidium and cesium react vigorously or even explosively with cold water. In each case, the aqueous metal hydroxide and hydrogen gas are produced, as shown:

\[ 2X_{(s)} + 2H_2O_{(l)} \rightarrow 2XOH_{(aq)} + H_{2(g)}\]

where \(X\) is any Group 1 metal. In each of the following descriptions, a very small portion of the metal is dropped into a large container of water.

Details for the individual metals

  • Lithium: Lithium's density is only about half that of water, so it floats on the surface, fizzing and giving off hydrogen gas. It gradually reacts and disappears, forming a colorless solution of lithium hydroxide. The reaction generates heat slowly, and lithium's melting point is too high for it to melt (this is not the case for sodium).

  • Sodium: Sodium also floats in water, but enough heat is given off to melt the sodium (sodium has a lower melting point than lithium) and it melts almost at once to form a small silvery ball that moves rapidly across the surface. The ball leaves a white trail of sodium hydroxide, which soon dissolves to give a colorless solution of sodium hydroxide.

    The sodium moves because it is pushed by the hydrogen produced during the reaction. If the sodium becomes trapped on the side of the container, the hydrogen may catch fire and burn with an orange flame. The color is due to contamination of the normally blue hydrogen flame with sodium compounds.

  • Potassium: Potassium behaves like sodium except that the reaction is faster and enough heat is given off to ignite the hydrogen. This time the hydrogen flame is contaminated by potassium compounds, so the flame is lilac-colored.
  • Rubidium: Rubidium sinks because it is less dense than water. It reacts violently and immediately, with everything leaving the container. Rubidium hydroxide solution and hydrogen are formed.
  • Cesium: Cesium explodes on contact with water, possibly shattering the container. Cesium hydroxide and hydrogen are formed.

Note: Summary of the trend in reactivity

The Group 1 metals become more reactive towards water down the group.

The Net Enthalpy Changes (Thermodynamics)

It is tempting to conclude that because the reactions get more dramatic down the group, the amount of heat given off increases from lithium to cesium. This is not the case. The table below gives estimates of the enthalpy change for each of the elements undergoing the reaction with water:

\[ X (s) + H_2O(l) \rightarrow XOH(aq) + \dfrac{1}{2} H_2 (g) \]

 

  \(\Delta H\) (kJ / mol)
Li -222
Na -184
K -196
Rb -195
Cs -203

 

There is no consistent pattern in these values; they are all very similar, and counter intuitively, lithium releases the most heat during the reaction. The differences between the reactions are determined at the atomic level. In each case, metal ions in a solid are solvated, as in the reaction below:

\[ X(s) \rightarrow X^+(aq) + e^-\]

The net enthalpy change for this process can be determined using Hess's Law, and breaking it into several theoretical steps with known enthalpy changes.

\[ X(s) \rightarrow X(g)\]

\[ X(g) \rightarrow X^+(g) + e^-\]

  • The final enthalpy change is the hydration enthalpy, or the heat released when the gaseous ion comes into contact with water.

\[ X^+(g) \rightarrow X^+(aq)\]

These values are tabulated below (all energy values are given in kJ / mol):

  atomization energy 1st IE hydration enthalpy total
Li +161 +519 -519 +161
Na +109 +494 -406 +197
K +90 +418 -322 +186
Rb +86 +402 -301 +187
Cs +79 +376 -276 +179

There is no overall trend in the overall reaction enthalpy, but each of the component input enthalpies (in which energy must be supplied) decreases down the group, while the hydration enthalpies increase:

  1. The atomization energy is a measure of the strength of the metallic bond in each element. This decreases as the size of the atoms and the length of the metallic bond increase. The delocalized electrons are further from the attraction of the nuclei in the larger atoms.
  2. The first ionization energy decreases because the electron being removed is more distant from the nucleus with each progressive atom. The extra protons in the nucleus are screened by additional layers of electrons.
  3. The hydration enthalpy is a measure of the attraction between the metal ions and lone pairs on water molecules. As the ions increase in size, the water molecules are farther from the attraction of the nucleus. The extra protons in the nucleus are again screened by the extra layers of electrons.

The summation of these effects eliminates any overall pattern. Knowing the atomization energy, the first ionization energy, and the hydration enthalpy, however, reveals useful patterns.

Activation Energies (Kinetics)

Consider the energy input terms:

  atomization energy 1st IE total
Li +161 +519 +680
Na +109 +494 +603
K +90 +418 +508
Rb +86 +402 +488
Cs +79 +376 +455

 

A steady decrease down the group is apparent. From lithium to cesium, less energy is required to form a positive ion. This energy will be recovered (and overcompensated) later, but must be initially supplied. This process is related to the activation energy of the reaction.

Note

The lower the activation energy, the faster the reaction.

Although lithium releases the most heat during the reaction, it does so relatively slowly—not in one short, sharp burst. Cesium, on the other hand, has a significantly lower activation energy, and so although it does not release as much heat overall, it does so extremely quickly, causing an explosion.

Explaining the increase in reactivity down the group

The reactions proceed faster as the energy needed to form positive ions falls. This is in part due to a decrease in ionization energy down the group, and in part to a decrease in atomization energy reflecting weaker metallic bonds from lithium to cesium. This leads to lower activation energies, and therefore faster reactions.

Contributors

Jim Clark (Chemguide.co.uk)