Structure and Bonding

What is a Bond?

You are often asked to define a bond in exams. There are many kinds of bonds but all of their definitions answer these 2 basic questions: 

What are the particles? 

Why do they stick together? (Hint: it’s always electrostatic attractions! You just need to describe them.)

Ionic bond: The electrostatic attraction between oppositely charged ions. 

Covalent bond: The electrostatic attraction of two nuclei to their shared pair of electrons.

Metallic bond: The electrostatic attraction between metal cations and their delocalised electrons. 

Sigma bond: A covalent bond formed from a direct or 'head-on' combination of two p orbitals, or s with p, or s with s.

Pi bond: A covalent bond formed from the 'side-on' overlap between p orbitals (or d orbitals), with its electron density found above and below the bond axis. 

Hydrogen bond: A type of dipole-dipole intermolecular attraction. A ∂+ hydrogen atom, bonded to an electronegative atom in one molecule, is attracted to the  ∂-  atom that is bonded to a hydrogen in another molecule.

London dispersion forces (LDFs): The electrostatic attraction between a temporary, instantaneous dipole on one molecule and the induced dipole on another molecule. LDFs are present in everything except ideal gases.

Van der Waals forces: All intermolecular forces except hydrogen bonds (according to IB)*.

* Hydrogen bonds are a type of dipole-dipole interaction, and dipole-dipole interactions are VDW forces, so H-bonds are in fact VDW forces. However, H-bonds are stronger than most other VDW forces, so the IBDP counts them as a different type.🙃

Structure and Properties

If you are asked to explain a property, you must tell us how the material behaves according to its structure - simply stating the type of bonding alone does not usually explain why something dissolves or conducts electricity or decomposes etc. 

Structure refers to the type of particles that the substance is composed of, where they are held, how they move, if there are any repeating units (a lattice) or if it consists of discrete molecules, etc. If the substance is molecular, you should mention the polarity of covalent bonds (meaning any ∂+ or ∂-  sites), and remember to check if the molecule has any resonance structures, which occur when there is more than 1 possible place to put a double bond (so pi electrons are delocalised over part or all of the molecule). 

Physical properties describe how a material looks, feels or behaves. Examples include colour, conductivity, melting and boiling points, viscosity, volatility and density. Chemical properties describe the material’s role in a chemical reaction, such as oxidizing/reducing agent, acid/base, nucleophile/electrophile, catalyst. You may also describe the vigour of a reaction; whether the material is stable or reactive in given conditions. 

Models and Theories of Bonding

In your IBDP course you will build on your knowledge of the various models of bonding (ionic, covalent, metallic) but you will also begin to see differences in the theories of bonding used to explain how bonds form, especially in molecules. Lewis theory, which describes the sharing of electrons to achieve a full outer shell, cannot predict the shape of molecules or the strength of bonds. VSEPR theory allows us to determine the shape of a molecule but not the energy of a bond. 

Valence Bond theory builds on these ideas to consider the symmetry of the atomic orbitals. The regular s and p orbitals exist mostly along the axes and do not overlap well with other atoms' orbitals if they are approaching between the axes to make eg. a tetrahedron. However, if we mix or hybridise these orbitals to form sp, sp2 or sp3 orbitals, we can explain the shape of a molecule. 

For the sp2 hybridised carbon (like in ethene or methanal), we see that one of the atomic p orbitals is left unhybridised so that it can form a pi bond with the neighbouring atom. This means that the hybridised orbitals must all exist on the same plane, as there are only two dimensions allowed by the mixing of one s and two p orbitals. VSEPR then tells us that the electron domain geometry is trigonal planar. Carbon dioxide molecules are linear; the carbon is sp hybridised. It is 1-dimensional because the two hybridised orbitals only have s and px character, no py or pz character. 

To explain expanded octets we must invoke the d orbitals, so the valence shell of sulphur in the octahedral SF6 molecule is sp3d2 hybridised. Elements in period 1 or 2 cannot use their d orbitals for hybridisation or any kind of bonding, as the lowest energy 3d orbitals are still much more energetic than the highest energy occupied 2p orbitals, so electrons are not promoted this far.  Here is an interactive to help you understand the geometry of s, p, d and hybrid orbitals - try grabbing the molecule and moving it around.

A more comprehensive model for bonding is Molecular Orbital Theory. MO theory is not found in any courses at this level and you don't need to know about it. It considers molecular orbitals as extending over the molecule as a whole and gives more accurate predictions of the energies of bonding and antibonding molecular orbitals, so the stability of a molecule can be assessed at a glance from an MO diagram. VB and MO theories are not necessarily in competition with each other; there are some questions that are better answered by one model or the other, but the models themselves do not disagree. Broadly speaking, you could say that VB theory is more concerned with the valence electrons’ locations and the shapes of molecules. MO theory is more concerned with the energies of molecular orbitals and the mathematical study of their symmetry according to group theory (this is why you haven’t met it yet and should not worry about it!)

Try the ‘common mistakes’ questions below to see if you can avoid all the pitfalls…

1. Explain how ionic compounds only conduct electricity when molten or dissolved.

2. Describe the structure of lithium iodide (avoid saying ‘molecule’ or ‘intermolecular’).

3. Explain why metals are malleable and ionic compounds are brittle. (This is two questions!)

4. Explain the trends in melting points for Group 1 and Group 7 elements. (Again, treat this as two questions.)

5. How could you use a bimetallic strip to make a thermostat? (Thermal expansion - particles don’t change size!)

6. Explain the trends in viscosity and boiling point of the fractions in crude oil. (Surface area/shape of molecules and LDFs)