Evidence of Covalent Bonding (Edexcel International AS Chemistry): Revision Note

Evidence of Covalent Bonding

• Covalent bonding occurs between two non-metals

• A covalent bond involves the electrostatic attraction between nuclei of two atoms and the electrons of their outer shells

• No electrons are transferred but only shared in this type of bonding

• When a covalent bond is formed, two atomic orbitals overlap and a molecular orbital is formed

• Covalent bonding happens because the electrons are more stable when attracted to two nuclei than when attracted to only one

The positive nucleus of each atom has an attraction for the bonding electrons shared in the covalent bond

• In a normal covalent bond, each atom provide one of the electrons in the bond. A covalent bond is represented by a short straight line between the two atoms, H-H

• Covalent bonds should not be regarded as shared electron pairs in a fixed position; the electrons are in a state of constant motion and are best regarded as charge clouds

A representation of electron charge clouds. The electrons can be found anywhere in the charge clouds

Bond Polarity

Non-metals are able to share pairs of electrons to form different types of covalent bonds Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas This makes each atom more stable

• When two atoms in a covalent bond have the same electronegativity the covalent bond is nonpolar

The two chlorine atoms have identical electronegativities so the bonding electrons are shared equally between the two atoms

When two atoms in a covalent bond have different electronegativities the covalent bond is polar and the electrons will be drawn towards the more electronegative atom As a result of this: The negative charge centre and positive charge centre do not coincide with each other This means that the electron distribution is asymmetric The less electronegative atom gets a partial charge of δ+ (delta positive) The more electronegative atom gets a partial charge of δ- (delta negative)

• The greater the difference in electronegativity the more polar the bond becomes until the bond becomes ionic

Cl has a greater electronegativity than H causing the electrons to be more attracted towards the Cl atom which becomes delta negative and the H delta positive

The octet rule

• In some instances, the central atom of a covalently bonded molecule can accommodate more or less than 8 electrons in its outer shell

  • Being able to accommodate more than 8 electrons in the outer shell is known as ‘expanding the octet rule’

    • Atoms from Period 3 and below can accommodate more than 8 electrons as the d-orbitals are accessible

    • Common examples of elements that can have an 'expanded octet' are phosphorous and sulfur

    • For example, PCl₅, which has 10 electrons around the central phosphorus atom

    • In SF₆, there are 12 electrons around the central sulfur atom

  • Accommodating less than 8 electrons in the outer shell means that the central atom is ‘electron deficient’

    • Boron has 3 electrons in its outer shell, 2s²2p¹

    • When it forms a covalent compound, these three electrons are paired

    • For example, BF₃ which has 6 electrons around the central boron atom

Covalent bonding & simple covalent lattice structures

• Covalent bonding can be responsible for substances that have many different structures and therefore different physical properties

• Small molecules such as H₂O and N₂ are simple units made from covalently bonded atoms

• These simple molecules contain fixed numbers of atoms

• Simple covalent lattices have low melting and boiling points

  • These compounds have weak intermolecular forces between the molecules

  • Only little energy is required to break the lattice

• Most compounds are insoluble with water

  • Unless they are polar and can form hydrogen bonds (such as sucrose)

• They do not conduct electricity in the solid or liquid state as there are no charged particles

  • Some simple covalent compounds do conduct electricity in solution, but this is because their interaction with water produces ions, for example, HCl which forms H⁺ and Cl^- ions when in aqueous solution 

• Buckminsterfullere, C₆₀, is an exception to some of these general points about simple molecules 

  • Buckminsterfullerene and other fullerenes are spherical networks of carbon atoms

  • They are made up of large molecules but they are not classed as giant covalent structures

  • Compared to other simple covalent molecules, buckminsterfullerene has a higher melting and boiling point

    • Fullerenes are generally larger molecules so will have larger intermolecular forces because of the number of electrons within the molecule which require more energy to overcome

  • Fullerenes are good insulators as, despite having delocalised electrons within their structure, these cannot pass between molecules and therefore cannot conduct electricity 

Covalent bonding & giant covalent lattice structures

• Giant covalent structures have a huge number of non-metal atoms bonded to other non-metal atoms via strong covalent bonds

• These structures can also be called giant lattices and have a fixed ratio of atoms in the overall structure

• Some of the common macromolecules you should know about include diamond and graphite 

• Giant covalent lattices have very high melting and boiling points

  • These compounds have a large number of covalent bonds linking the whole structure

  • A lot of energy is required to break the lattice

• The compounds can be hard or soft

  • Graphite is soft as the forces between the carbon layers are weak

  • Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds

• Most compounds are insoluble with water

• Most compounds do not conduct electricity however some do

  • In graphite, each carbon atom has three outer electrons that are involved in bonding to other carbon atoms, leaving one electron which becomes delocalised between the carbon layers and can move along the layers when a voltage is applied

  • Diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom are involved in a covalent bond so there are no freely moving electrons available