Examples of covalent elements/compound examples described: hydrogen H₂, chlorine Cl₂, hydrogen chloride HCl, water H₂O, ammonia NH₃, methane CH₄, oxygen O₂, carbon dioxide CO₂, ethene C₂H₄, nitrogen N₂, ethane C₂H₆, chloromethane CH₃Cl, methanol CH₃OH

More advanced level chemistry notes on the shapes and bond angles of molecules and ions

Covalent bonds are formed by atoms sharing electrons to form molecules. This type of bond usually formed between two non-metallic elements. The molecules might be that of an element i.e. one type of atom only OR from different elements chemically combined to form a compound.

This kind of bond or electronic linkage does act in a particular direction i.e. along the 'line' between the two nuclei of the atoms bonded together, this is why molecules have a particular shape. In the case of ionic or metallic bonding, the electrical attractive forces act in all directions around the particles involved.

Example 1: two hydrogen atoms (1) form the molecule of the element hydrogen H₂

 and combine to form where both atoms have a pseudo helium structure of 2 outer electrons around each atom's nucleus. Any covalent bond is formed from the mutual attraction of two positive nuclei and negative electrons between them. The nuclei would be a tiny dot in the middle of where the H symbols are drawn! H valency is 1.

Example 2: two chlorine atoms (2.8.7) form the molecule of the element chlorine Cl₂

 and combine to form where both atoms have a pseudo argon structure of 8 outer electrons around each atom. All the other halogens would be similar e.g. F₂, Br₂ and I_{2 etc}. Valency of halogens like chlorine is 1 here.

Example 3: one atom of hydrogen (1) combines with one atom of chlorine (2.8.7) to form the molecule of the compound hydrogen chloride HCl

 and combine to form where hydrogen is electronically like helium and chlorine like argon. All the other hydrogen halides will be similar e.g. hydrogen fluoride HF, hydrogen bromide HBr and hydrogen iodide HI.

Note: Hydrogen chloride gas is a true covalent substance consisting of small HCl molecules. If the gas is dissolved in a hydrocarbon solvent like hexane or methylbenzene it remains as HCl molecules and because there are no ions present, the solution does not conduct electricity. However, if hydrogen chloride gas is dissolved in water, things are very different and the HCl molecules split into ions. Hydrochloric acid is formed which consists of a solution of hydrogen ions (H⁺) and chloride ions (Cl^-). The solution then conducts electricity and passage of a d.c. current causes electrolysis to take place forming hydrogen and chlorine.

Reminder: How to work out formula of covalent compounds without going through some demanding electronic thinking is described on the "Elements, Compounds and Mixtures" page and it is followed by a section on naming compounds.

Example 4: two atoms of hydrogen (1) combine with one atom of oxygen (2.6) to form the molecule of the compound water H₂O

 and and combine to form so that the hydrogen atoms are electronically like helium and the oxygen atom becomes like neon. The molecule can be shown as with two hydrogen - oxygen single covalent bonds (AS note: called a V or bent shape, the H-O-H bond angle is 105^o). Hydrogen sulphide will be similar, since sulphur (2.8.6) is in the same Group 6 as oxygen. Valency of oxygen and sulphur is 2 here.

Example 5: three atoms of hydrogen (1) combine with one atom of nitrogen (2.5) to form the molecule of the compound ammonia NH₃

three of  and one combine to form so that the hydrogen atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule can be shown as with three nitrogen - hydrogen single covalent bonds (AS note: called a trigonal pyramid shape, the H-N-H bond angle is 107^o). PH₃ will be similar since phosphorus (2.8.5) is in the same Group 5 as nitrogen. Valency of nitrogen or phosphorus is 3 here.

Example 6: four atoms of hydrogen (1) combine with one atom of carbon (2.4) to form the molecule of the compound methane CH₄

four of  and one of combine to form so that the hydrogen atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule can be shown as with four carbon - hydrogen single covalent bonds (AS note: called a tetrahedral shape, the H-C-H bond angle is 109^o). SiH₄ will be similar because silicon (2.8.4) is in the same group as carbon.

All the bonds in the above examples are single covalent bonds. Below are three examples 7-9, where there is a double bond in the molecule, in order that the atoms have stable Noble Gas outer electron arrangements around each atom. Carbon and silicon have a valency of 4.

More complex examples can be worked out e.g. involving C, H and O. In each case link in the atoms so that there are 2 around a H (electronically like He), or 8 around the C or O (electronically like Ne).

Example 7:  Two atoms of oxygen (2.6) combine to form the molecules of the element oxygen O₂. The molecule has one O=O double covalent bond . Oxygen valency 2.

Example 8:  One atom of carbon (2.4) combines with two atoms of oxygen (2.6) to form the compound carbon dioxide CO₂. The molecule can be shown as   with two carbon = oxygen double covalent bonds (AS note: called a linear shape, the O=C=O bond angle is 180^o). Valencies of C and O are 4 and 2 respectively.

Example 9:  Two atoms of carbon (2.4) combine with four atoms of hydrogen (1) to form ethene C₂H₄. The molecule can be shown as  with one carbon = carbon double bond and four carbon - hydrogen single covalent bonds (called a planar shape, its completely flat!, the H-C=C and H-C-H bond angles are 120^o). The valency of carbon is still 4.

Examples 10-13: The scribbles below illustrate some more complex examples. Can you deduce them for yourself? Ex. 10 nitrogen N₂; Ex. 11 ethane C₂H₆; Ex. 12 chloromethane CH₃Cl and Ex. 13 methanol CH₃OH. Electronic origin of the diagrams showing the outer electrons of N, C, Cl and O: N at. no. 7 (2.5), H at. no. (1), C at. no. 6 (2.4), Cl at. no. 17 (2.8.7) and O at. no. 8 (2.6) plus a variety of crosses and blobs! The valencies or combining power in theses examples are N 3, H 1, C 4, Cl 1 and O 2. From these you can work out others e.g. Ex. 12 can be used to derive the ox diagram for tetrachloromethane CCl₄.

• The electrical forces of attraction, that is the chemical bond, between atoms in a molecule are usually very strong, so,  most covalent molecules do not change chemically on moderate heating.

  • e.g. although a covalent molecule like iodine, I₂, is readily vapourised on heating, it does NOT break up into iodine atoms I. The I-I covalent bond is strong enough to withstand the heating and the purple vapour still consists of the same I₂ molecules as the dark coloured solid is made up of.

• So why the ease of vaporisation on heating?

  • The electrical attractive forces between individual molecules are weak, so the bulk material is not very strong physically and there are also consequences for the melting and boiling points.

• These weak electrical attractions are known as intermolecular forces and are readily weakened further on heating. The effect of absorbing heat energy results in increased the thermal vibration of the molecules which weakens the intermolecular forces. In liquids the increase in the average particle kinetic energy makes it easier for molecules to overcome the intermolecular forces and change into a gas or vapour.  Consequently, small covalent molecules tend to be volatile liquids with low boiling points, so easily vapourised, or low melting point solids.

  • On heating the inter-molecular forces are easily overcome with the increased kinetic energy of the particles giving the material a low melting or boiling point and a relatively small amount of energy is needed to effect these state changes.

  • Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.

  • This contrasts with the high melting points of giant covalent structures with their strong 3D network.

  • Note: The weak electrical attractive forces between molecules, the so called intermolecular forces should be clearly distinguished between the strong covalent bonding between atoms in molecules (small or giant), and these are sometimes referred to as intramolecular forces (i.e. internal to the molecule).

• Covalent structures are usually poor conductors of electricity because there are no free electrons or ions in any state to carry electric charge.

• Most small molecules will dissolve in some solvent to form a solution.

  • This again contrasts with giant covalent structures where the strong bond network stops solvent molecules interacting with the particles making up the material.

• The properties of these simple small molecules should be compared and contrasted with those molecules of a giant covalent nature (next section).

  • Apart from points on the strong bonds between the atoms in the molecule and the lack of electrical conduction, all the other properties are significantly different!

• More advanced notes on intermolecular forces

• More advanced level chemistry notes on the shapes and bond angles of molecules and ions