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Unusual properties of water

Water can assume various physical states within a range of 100 degrees.

After introductory pictures of different forms of water (for example as rain, haze, frost, snow or ice), the water is first compared with nitrogen and oxygen.

Although the molecules of these three substances are roughly the same size, nitrogen (N.2) as well as oxygen (O2) - at comparable temperatures - only gaseous, water (H.2O), on the other hand, in all three states of aggregation. A space-filling model of the water molecule serves to explain this phenomenon.

Based on the electron configuration, the notation for the oxygen atom - with dots for the unpaired and with dashes for the paired electrons - is clarified. These explanations supplement the electron configuration of the two hydrogen atoms as well as a trick illustration of the overlap of the oxygen and hydrogen forbitals to the Molecular orbital of the water.

Electrical deflection of a jet of water

An experiment follows in which a jet of water is deflected by an electrically charged plastic rod. This is the result of the different electronegativity of oxygen and hydrogen, through which the binding electrons are more strongly attracted to the oxygen.

Different electronegativity of elements of the PSE - please click on the magnifying glass.

An overview of the periodic table with the values ​​for the electronegativity leads finally to a molecular model of water, which shows the different electronegativity of oxygen and hydrogen.

Dipole character of water

Model experiment on the dipole character of water molecules

The different electronegativity of oxygen and hydrogen causes an uneven charge distribution in the water molecule. This is expressed in the so-called Dipole character of the water.

Evaporation of liquid oxygen at room temperature

Compared to oxygen and nitrogen, which due to the weak van der Waals forces only change from a gaseous to a liquid state at temperatures below -183 ° C (or -196 ° C), this case occurs with water because of the additional strong dipole forces already at room temperature. An experiment in which poured out liquid oxygen evaporates immediately can be taken as evidence of this fact.

Structure of ice crystal cells due to the hydrogen bond

An ice crystal demonstrates how Hydrogen bonds act as relatively weak electrical forces between neighboring water molecules.

Hydrogen bond

When heated, however, the ordered crystal structure of ice breaks down to form denser liquid water. The fact that ice floats in water proves that the volume of liquid water is actually less than - with the same mass - of solid water.

Four compounds with water-like molecules

Hydrogen sulfide (H.2S), hydrogen selenide (H.2Se) and hydrogen telluride (H.2Te) are very similar to water in structure and in the size of their molecules. Using the example of hydrogen sulfide (H.2S) explains a crucial difference to water:

Different electronegativity of oxygen and sulfur - please click on the magnifying glass.

Water is liquid at room temperature, whereas hydrogen sulfide is gaseous. A comparison of the electronegativity of oxygen, hydrogen and sulfur makes it clear that the difference in electronegativity (ΔEN) in the water molecule is greater than in the hydrogen sulfide molecule. H2S is therefore a weaker dipole than water. Its molecules are therefore less attracted to each other, which means that the compound remains in a gaseous state at room temperature.

Different electronegativity difference between water and hydrogen sulfide

Comparing the boiling temperatures of H.2O, (H2S), (H2Se) and (H2Te), then, as expected, they decrease from H2O to (H2S) by around 160 ° C, but then rise again to –2 ° C, although the ΔEN decreases to 0.

Increasing boiling temperature due to increasing Van der Waals forces - please click on the magnifying glass.

This decrease in polarity is overcompensated for by the increase in the van der Waals forces in the larger molecules.

In the experiment, a charged plastic rod directs a liquid jet of carbon tetrachloride (CCl4) Not although this should be the case due to the difference in electronegativities (ΔEN = 0.5) in the molecule. The reason: The effect of the existing, polar bonds is canceled out by a molecular symmetry that is established. Despite polar bonds, the molecule as a whole is not a dipole.

At the end of the broadcast, the bond angle of 105 ° in the water molecule is discussed. This can be explained by the fact that in the case of oxygen, the electrons of the 2nd shell are equal in terms of energy. This creates four similar 2sp3-Hybrid orbitals, which, however, are occupied differently with electrons. Since two of these orbitals contain only one electron, they are suitable for bonding with hydrogen atoms. With this connection, different orbital shapes are created, which cause the bond angle of 105 °.