The most significant intermolecular force acting between water (H2O) molecules is the hydrogen bond.
Weaker intermolecular forces act between other Group 16 hydride molecules.
The melting point and boiling point of water (H2O) molecules is unexpectedly high due to the stronger hydrogen bonds acting between water molecules.
Structure of Group 16 Hydrides
Consider the structure of a hydrogen, H, atom.
Hydrogen has atomic number, Z, 1 so it has only one positively charged proton in the nucleus of the aotm.
A neutral hydrogen atom must therefore have just one negatively charged electron occupying the space around the nucleus in order for the charge to be balanced.
The Lewis Structure (electron dot diagram) for an atom of hydrogen is shown on the right.
An atom of hydrogen can achieve the stable electron configuration of helium, He, by contributing its 1 electron to be shared between itself and another atom, such as an atom of a Group 16 element.
So, let's consider the structure of each Group 16 element:
Group 16 Elements
Atomic Number (Z)
No. Valence Electrons
Atomic number (Z), the number of protons in the nucleus of the atom and therefore the number of electrons present in an atom of the element, increases down the group from oxygen (O) to livermorium (Lv).
The number of valence electrons, the number of electrons in the highest energy level, remains constant at 6 for the atoms of Group 16 elements.
All Group 16 elements have 6 valence electrons.
The Lewis Structure, electron dot diagram, for any Group 16 element, represented by the letter E, is shown on the right.
In order to achieve a stable electron configuration with 8 electrons in the valence shell (highest energy level), each atom of a Group 16 element can share 1 valence electron with each of 2 hydrogen atoms.
That is, each Group 16 element forms a covalent bond with 2 hydrogen atoms.
So all the hydrides of Group 16 elements are covalent compounds with the general molecular formula H2E.
The Lewis Structures (electron dot diagram) and Valence Structures for each of the common Group 16 hydrides is shown below:
hydrogen sulfide H2S
hydrogen selenide H2Se
hydrogen telluride H2Te
Shape of the Group 16 Hydride Molecules
The shape of the Group 16 hydride molecules is determined by the number of bonding pairs of electrons and lone pairs of electrons (non-bonding pairs) around the central Group 16 atom.
The repulsion between the lone pairs of electrons forces these pairs as far away as possible from each other.
If an atom of element E had made 4 covalent bonds to atoms of another element using 4 bonding pairs of electrons, we would have a molecule with the molecular formula EX4 and all the bonding pairs of electrons would repel each other equally resulting in a symmetrical tetrahedral arrangement of bonded atoms with the angles between each bonding pair (bond angles) all 109.5° as shown on the right:
If element E makes only 3 covalent bonds to element X, EX3, there are now 3 bonding pairs of electrons and 1 lone pair (non-bonding pair) of electrons.
This introduces an asymmetry into the molecule.
The lone pair - bonding pair repulsion is greater than the bonding pair - bonding pair repulsion, so the angle between the lone pair and bonding pair of electrons will now be greater than 109.5°, while the angle between the bonding pairs will be less than 109.5°, that is, we now have a distorted tetrahedral arrangement of electron pairs around the central atom E as shown on the right:
When Element E forms 2 covalent bonds with an atom such as hydrogen, 2 pairs of electrons are bonding pairs and 2 pairs of electrons are lone pairs (non-bonding pairs). The repulsion between the lone pairs of electrons is greater than the repulsion between lone pairs and the bonding pairs, which is greater than the repulsion between the bonding pairs of electrons, so the bonding pairs of electrons are pushed even closer together.
H-E-H bond angle
H-O-H bond angle a bit less than tetrahedral bond angle (109.5°)
These H-E-H bond angles are significantly less than tetrahedral bond angle, and are all about 90°
Question: What shape are all Group 16 hydride molecules?
In general, Group 16 atoms have a greater ability to attract electrons towards themselves than hydrogen atoms.
This property is referred to as electronegativity.
The electronegativity of hydrogen and the common Group 16 elements is shown in the table below:
Notice that oxygen is significantly more electronegative than the other Group 16 atoms.
Note the similar values for H and Te
Since the electronegativity of the Group 16 atoms decreases down the group, the difference between the electronegativity of the hydrogen atom and the Group 16 atom also decreases down the group:
Group 16 atom
Electronegativity of Group 16 atom
Electronegativity of H atom
Difference in electronegativity
3.5 - 2.1 = 1.4
O-H bond is significantly more polar.
2.5 - 2.1 = 0.4
S-H and Se-H bonds have similar polarity and are significantly less polar than the O-H bond
2.4 - 2.1 = 0.3
The difference in the electronegativities of Group 16 elements and hydrogen results in a polar covalent bond, that is, compared to the hydrogen atoms the Group 16 atom generally has a greater attraction for the shared electrons so the electrons of the shared electron pair spend more time around the nucleus of the Group 16 atom resulting in a very, very slight (partial) negative charge on the Group 16 atom (δ-) and a partial positive charge (δ+) on the hydrogen atoms.
Since O and H have the greatest difference in electronegativity, the partial negative charge on oxygen in the O-H bond will be much greater than the partial negative charge on the sulfur in the S-H bond but this is then only very slightly greater than the partial negative charge on the selenium atom in the Se-H bond.
We saw above that due to the presence of 2 lone pairs of electrons (non-bonding pairs) on the central Group 16 atom, the shape of the Group 16 hydride molecule is NOT linear, it is bent and therefore not symmetrical.
In a solid, the molecules of Group 16 hydrides are held together by attractive intermolecular forces resulting in a 3-dimensional lattice.
In order to melt the solid, energy must be supplied to weaken these attractive intermolecular forces so that the molecules can "roll over" each other in the liquid state.
General Word Equation
General Chemical Equation (E = Group 16 element)
The reaction is reversible, that is, we could remove energy from the system by cooling it in order to change the liquid back into a solid and this is why we use the equilibrium arrow ( ⇋ ) rather then a single directional arrow ( → ) in the chemical equation above.
Now, consider the melting points of Group 16 hydrides as shown in the table below:
Relative Molecular Mass1
Melting Point / °C
Melting point of water anomalously high!
Gradual increase in melting points down the rest of the group as the relative molecular mass of the molecules increases.
A graph of these melting points against relative molecular mass is shown below:
The anomalously high melting point of water (relative molecular mass 18) becomes readily apparent from the graph!
This means that it requires a lot more energy to melt solid water (ice) than it does to melt any other Group 16 hydride.
For this to be true, the intermolecular forces holding the water molecules together in a lattice must be much stronger than the intermolecular forces holding other Group 16 hydride molecules together in their lattices.
And we have seen above that the water molecule is significantly more polar than the other Group 16 hydride molecules so the attraction between Hδ+ and Oδ- between water molecules is stronger than the attraction of Hδ+ and Eδ- between other Group 16 hydride molecules.
This stronger intermolecule force is called a hydrogen bond and acts between water molecules.
In the diagram below, very strong covalent bonds between O and H in a single water molecule are shown as black solid lines.
The hydrogen bonds, which are much weaker than covalent bonds, act between the water molecules and are shown as red dotted lines in the diagram below.
Melting H2O(s) requires more energy to weaken the stronger hydrogen bonds between water molecules in the lattice compared to the amount of energy required to weaken the weaker intermolecular forces holding the other Group 16 hydride molecules together in a lattice.
Therefore the melting point of H2O(s) is greater than the melting points of the other Group 16 hydrides.
Question: Describe the general trend for melting points of Group 16 hydrides.
In a liquid, the molecules of Group 16 hydrides are attracted to each other by intermolecular forces.
In order to boil the liquid, energy must be supplied to completely overcome these attractive intermolecular forces so that the molecules can escape from the liquid phase and exist independently in the gaseous phase.
General Word Equation
General Chemical Equation (E = Group 16 element)
The reaction is reversible, that is, we could remove energy from the system by cooling it in order to change the gas back into a liquid and this is why we use the equilibrium arrow (⇋) rather then a single directional arrow ( → ) in the chemical equation above.
Now, consider the boiling points of Group 16 hydrides as shown in the table below:
Relative Molecular Mass2
Boiling Point / °C
Boiling point of water is anomalously high!
Gradual increase in boiling points down the rest of the group as the relative molecular mass of the molecules increases.
The anomalously high boiling point of water (relative molecular mass 18) becomes readily apparent from the graph!
This means that it requires a lot more energy to boil liquid water than it does to boil any other Group 16 hydride.
Much more energy is required to overcome the stronger hydrogen bonds that act between water molecules in liquid water compared to the lesser amount of energy needed to overcome the weaker intermolecular forces holding the other Group 16 hydride molecules together in the liquid phase.
Therefore, the boiling point of water is much higher than the boiling point of the other Group 16 hydrides.
Question: Water is considered to have unique properties. Why?