Introduction to bonding

Introduction to bonding


In this lesson, we will learn:
  • The two main categories of bonding in atoms and molecules.
  • How different bonds and their varying strengths leads to a variety of chemical properties.
  • Why different chemical substances display different types of bonding.
  • The link between the bonding, structure and properties of chemical substances.


  • There are over 100 different elements in the periodic table but millions of different chemical 'species' – chemicals unique due to their particular arrangement of atoms. This is possible because there are so many different combinations of elements and how they can join together. The different ways that atoms and molecules combine to form larger or more organized structures, and the attractive forces that make it happen is known as bonding.

  • Bonding is a very general word – it can be used to describe any of the attractive forces that act between or inside molecules, and if you are asked to 'describe the bonding' in a substance you should talk about any of the attractive forces present. Before learning any of the particular types of bonding, recall the principles of electrostatic forces we saw in the Periodic Table and Elements chapter. Knowing them well will help explain why different types of bonding exist:
    • #1: Oppositely charged particles attract each other, while particles of like charge repel each other.
    • #2: The greater the charge difference of two particles, the greater their force of attraction (for example, the attractive force of a 2+ charge attracting a 2- charge is greater than the attractive force of a 1+ charge attracting a 1- charge).
    • #3: Attractive forces between oppositely charge particles decrease with distance.
    • #4: Repulsive forces between like charged particles decrease with distance.

  • There are many different types of bond or attractive force but we can put them into two broad categories:
    • Forces that hold the atoms of a molecule or compound together, acting between the atoms inside each of the molecules, are intramolecular forces. All chemical bonds are intramolecular forces.
    • Forces and interactions in between molecules are called intermolecular forces.

  • Intramolecular forces (ionic and covalent bonds) are much stronger than intermolecular forces. When a chemical substance melts or evaporates, it is the intermolecular forces being overcome and broken, not the intramolecular forces!

  • How an atom of an element bonds is largely determined by the valence of the atom – the number of unpaired outer shell electrons. The valence of an atom practically tells you how many bonds the atom can make:
    • In Electronic structure: 288 rule and Electronic structure: Subshells, we saw that the shape of the Periodic Table is made to show the different electron subshells or orbitals that atoms have – the s, p, d block etc.
    • When atoms bond with other atoms, these orbitals mix together. So to write 'valence electrons' in bonding, show the outer shell as four equal orbitals, not separate, different s and p sections. Fill in valence shells by adding single unpaired electrons to the four orbitals, then start pairing them up with the 5th onward.
    • Because paired electrons generally don't bond, only the unpaired electrons are considered available to make bonds.
    • Across the Periodic Table from groups 1-8 (ignoring the transition metals), the valence of the groups are: 1, 2, 3, 4, 3, 2, 1, 0. As you can see, the electron configuration is hugely important to how an atom can bond. Being comfortable with electron configurations will help your understanding of bonding a lot!

    • overview of the particle properties

  • The properties a chemical displays are due to the types of bonding and interactive forces between its atoms and molecules, and the types of bonding a chemical shows is because of its atoms' valence, and the electrons being able to make certain bonds in order to gain a full valence shell.
    • Example 1: Neon is a noble gas with 8 valence electrons all paired up. Therefore neon atoms have a valence of zero, and they don't make bonds between atoms to form molecules or compounds.
      • This means they don't have strong intramolecular forces; a sample of neon gas exists as millions of single atoms freely floating in space.
      • Therefore they are entirely limited to weak intermolecular forces and so have very low boiling points, are gases at room temperature and cannot conduct heat or electricity.
    • Example 2: Carbon atoms have four unpaired valence electrons, meaning carbon atoms have a valence of four and each atom can make four covalent bonds (a type of intramolecular force) with other atoms.
      • Because each atom can 'connect' with a strong bond to four others, pure carbon is found in some forms that have the atoms in a single giant structured network, made of millions of individual carbon atoms all covalently bonded together. This is the structure of diamond.
      • When doing this, because the structure is effectively one giant molecule and no one atom can be disturbed without disrupting the entire structure, diamond is the hardest substance known to man and, practically speaking, cannot be melted.

  • The valence bonding and interactive forces that exist in a chemical species gives rise to the structure of the molecules it makes, which dictate its properties. Understanding the link between bonding that leads to structure that leads to properties is crucial and will allow you to make some predictions about certain chemical substances even when given just the formula.
  • Introduction
    Introduction to bonding
    Why are there so many different chemicals?

    Types of "bonding" and forces.

    Working out the valence of an atom.

    Why is valence important for elements?

  • 1.
    Apply knowledge of electron shells to find the valence of elements.
    Show the valence electrons in an atom of:




  • 2.
    Understand the difference between types of bonding and forces.
    Explain whether covalent bonding is an intramolecular or intermolecular force.