Is H2S an ionic or a covalent compound?

The inorganic compound hydrogen sulphide (H2S) has the chemical formula H2S. If you’ve ever had the misfortune of performing a qualitative examination of sulfur-containing compounds, you’ll identify H2S as a gas that smells like rotten eggs (yuck!).

Is H2S an ionic or a covalent compound? H2S is a covalent molecule because the large size and larger charge on the sulphide anion (S2-) facilitate the development of a covalent bond, according to Fajan’s Rules. The sulphur and hydrogen atoms have a tiny electronegativity difference, which aids in the creation of a covalent bond.

Hydrogen and sulphur are the two elements that make up H2S. As a result, it’s a binary compound. Pure H2S is obtained in the industrial world by separating it from other natural gas constituent gases.

H2S appears to be a pretty simple molecule at first appearance. However, at very low temperatures ( 150 Kelvin) and very high pressures (> 100 GPa), this common compound is known to exhibit superconductivity, or zero resistance to electric current.

We perform a rigorous chemical study in this article to establish whether H2S is ionic or covalent.

What distinguishes a covalent bind from an ionic bond?

When two atoms share electrons to produce a stable electronic configuration, they form a covalent connection.

The two atoms in question gain and lose electrons to produce a negatively charged anion and a positively charged cation, respectively, in the case of ionic bonding.

The cation and anion acquire stability as a result of having a complete octet, and the cation-anion interaction boosts the molecule’s stability even more.

When the properties of their compounds are compared, the difference between the two can be further elucidated.

Ionic CompoundsCovalent Compounds
Usually crystalline solidsUsually fluids, seldomly solids
High Melting and Boiling Points (Very Strong Bonds)Low Melting and Boiling Points (Weak Bonds)
Soluble in Polar solvents but insoluble in Non-Polar solventsSoluble in Non-Polar solvents but insoluble in Polar solvents
High electrical conductivity in the dissolved or molten stateLow electrical conductivity

The conditions for the development of a covalent bond are as follows:

  1. Because ionisation potential is related to the energy necessary for ion production, both atoms should have a high ionisation potential.
  2. To aid in electron sharing, both atoms should have a strong electron affinity.
  3. The difference in electronegativity should be as small as feasible, because a greater contrast in electronegativity indicates more ionic nature and less electron sharing.

Fajan’s Rules and Polarization

In reality, neither an ionic nor a covalent chemical link is 100 percent. The degree of polarisation determines the percentage covalent quality of a chemical bond.

Polarization is defined as a cation’s tendency to distort an anion’s electron cloud and vice versa.

The degree of polarisation determines the covalent character of a chemical bond.

With a big positive charge, smaller cations produce a high degree of polarisation, while larger anions induce a strong negative charge.

Why is H2S a covalent compound?

Let’s look at the electron configuration of H and S to see what kind of bonding they have.

H (Atomic Number 1) electron configuration: 1s1

S (Atomic Number 16) electron configuration: [Ne] 3s2 3p4

H is one electron shy of obtaining Helium’s (Atomic Number 2) inert electron configuration, and S is two electrons short of achieving Argon’s inert electron configuration (Atomic Number 18).

Hydrogen and sulphur have values of 2.20 and 2.58, respectively, on Pauling’s electronegativity scale. The higher the value, the more electronegative the atom, and the lower the value, the more electropositive the atom.

Take the production of two H+ cations and one sulphide anion S2- as an example.

In terms of electron configuration, both ions are stable. However, the H+ cation is relatively tiny (a free proton), whereas the S2- anion is quite massive (170 pm). Covalent bonds between the three atoms must be established, according to Fajan’s Rules.

To visualise the sharing of electrons via covalent bonding, a Lewis structure of H2S must be drawn.

Lewis structure and VSEPR Theory are used to study H2S bonding.

To identify the structure of H2S, we use the following procedure:

Determine the centre atom in step one. We use S as the core atom in our example.

Step 2: Count the valence electrons in the molecule as a whole.

n1 = 6 (derived from S) + 2 x 1 (derived from H) = 8

Step 3: Determine how many electrons are required to complete the octet of all atoms.

8 x 1 + 2 x 2 = 12 n2 = 8 x (number of non-H atoms) Plus 2 x (number of H atoms)

Step 4: (n2 – n1)/2 = 4/2 = 2 = number of bond pairs

Step 5: Count how many non-bonding electrons there are.

8 – (12 – 8) = 4 n3 = n1 – (n2 – n1) = 8 – (12 – 8) = 4

Step 6: n3/2 = 4/2 = 2 = number of lone pairings

The Lewis structure of H2S is now ready to be drawn.

First, we place S in the centre and connect it to the H atoms with two bond pairs. The covalent bonds that bind S and H are formed by the two bond pairs.

It’s worth noting that the H atoms have acquired Ne’s electrical structure, but the S atom still lacks four electrons.

To correct this, the two lone pairs are placed on the S atom. The S atom has now acquired the Ar electron configuration, which is stable and inert.

Step 7: Determine all atoms’ formal charges.

The concept of formal charge can be used to assess the validity of our Lewis structure. An atom’s formal charge is defined as follows:

Valence Electrons – (0.5 x Bonding Electrons) – Non-Bonding Electrons = Formal Charge

The Lewis structure can be used to calculate the number of bonding and non-bonding electrons.

S atom formal charge = 6 – (0.5 x 4) – 4 = 6 – 6 = 0

Each H atom has a formal charge of 1 – (0.5 x 2) – 0 = – 0 = – 0 = – 0 = – 0 = – 0 =

0 + 0 + 0 = 0 Total Charge on the Molecule Equals Sum of Formal Charges on Atoms

This makes sense because H2S is a neutral, uncharged molecule. As a result, we have the correct Lewis structure.

VSEPR theory can be used to predict the molecule’s geometry after knowing the electron distribution in the molecule.

The molecule will bend due to the repulsion between the lone pairs on the S atom and the bond pairs on the H atoms.

H2S Characteristics

Chemical Characteristics

  1. The gas is poisonous when inhaled.
  2. Because S is less electronegative than O, H2S is a superior nucleophile in a polar aprotic solvent than H2O.
  3. It decomposes slowly in the air, forming elemental Sulfur.
  4. Because of hydrogen bonding and mild acidity, H2S is only marginally soluble in water.
  5. It forms metal sulphides when it reacts with metal ions.

Physical Characteristics

  1. Hydrogen sulphide (H2S) is a colourless gas with a distinct rotten egg stench.
  2. The tendency of H2S to persist in the gaseous phase is shown by its very low melting point (-82° Celsius or -116° Fahrenheit) and boiling point (-60° Celsius or -76° Fahrenheit).
  3. It’s a combustible gas.
  4. It has a somewhat higher density than air.
  5. By holding H2S under its own vapour pressure, it can be liquified.

H2S’s Applications

  1. Inorganic substances are analysed qualitatively using H2S. The existence of sulphide anions is frequently indicated by its presence. Bubbling the gas through the chemical solution can cause heavy metal ions like Pb(II), Cu(II), and Hg to precipitate (II).
  2. In metallurgy, H2S is injected to cause the valuable metal to precipitate as metal sulphides.
  3. Organosulfur compounds such as methanethiol, ethanethiol, and thioglycolic acid are synthesised using the nucleophilicity of H2S.
  4. Hydrogen sulphide is utilised to make sulphur alkali metal compounds, which are widely employed in the paper industry.


Because the S and H atoms share electrons, H2S is a covalent molecule. Because electrons are shared, all atoms have an inert electron configuration.

The development of two single bonds between the S and H atoms is due to two bond pairs of electrons. On the S atom, the remaining valence electrons appear as non-bonding electrons or lone pairs.

Read more: Molecular Geometry, Hybridization, and Polarity of SiS2 Lewis Structure

Misha Khatri
Misha Khatri is an emeritus professor in the University of Notre Dame's Department of Chemistry and Biochemistry. He graduated from Northern Illinois University with a BSc in Chemistry and Mathematics and a PhD in Physical Analytical Chemistry from the University of Utah.


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