Bonding and Structure

 
 

Ionic bonding

Atoms form ions when they gain or lose electrons to achieve a complete octet.

  • Group 1 and 2 metals tend to lose electrons and form positive ions.

  • Non-metals in groups 6 and 7 tend to gain electrons and form negative ions.

An ionic bond is the electrostatic attraction between positive and negative ions.

 
 

We can show how the metallic atom loses electrons and the non-metallic atom gains electrons by drawing a dot and cross diagram. Normally, we just draw the electrons present in the outer energy level and use different symbols for the electrons on each atom.

This dot and cross diagram shows how sodium — who has one outer electron — loses the electron to form a +1 ion. An atom of chlorine takes the electron from sodium, forming a -1 ion.

Ionic compounds can form giant lattices, with a large number of positive and negative ions forming numerous ionic bonds.

  • These structures have high melting and boiling points because ionic bonds are strong and require a great deal of energy to break.

  • This makes ionic compounds solid at room temperature.

  • They do not conduct electricity when solid but can conduct electricity when molten or as part of an aqueous solution. Remember that in order to conduct electricity, there need to be charged molecules that are free to move. The positive and negative ions can only move when the ionic compound is melted or dissolved in solution.

Covalent bonding

A covalent bond is formed when two atoms share electrons. When this happens there is an electrostatic attraction between the nuclei (which are positively charged) and the pair of shared electrons (which are negatively charged).

In the dot and cross diagram below, you can see how hydrogen forms a covalent bond by sharing its single electron with chlorine to form hydrogen chloride.

 
 

Molecules that consist of just a few atoms held together by covalent bonds are called simple covalent molecules. Examples include water, H2O, carbon dioxide, CO2 and ammonia, NH3.

Simple covalent molecules have low melting and boiling points because the strong covalent bonds remain intact when these molecules change state. Instead, it is the weak intermolecular forces that need to be broken when converting a simple covalent molecule between a solid, liquid and gaseous state.

Although all simple covalent molecules will have low melting and boiling points, those with a large molecular mass will have a higher melting and boiling point than a molecule with a smaller mass. This is because the weak intermolecular forces that are found between molecules form from electrons. Heavier molecules have a larger number of electrons which means they can form more intermolecular forces, thereby increasing their melting and boiling point.

If a compound forms double or triple bonds, you just need to draw two or three pairs of dots and crosses. The image below shows a dot-and-cross diagram for the double bonds in a molecule of carbon dioxide and a triple bond in a molecule of nitrogen.

 
 

Dative covalent bonds (aka coordinate covalent bonds) are a special type of covalent bond that is formed when both electrons come from the same species. For example, ammonia has a lone pair of electrons which it can use to form a dative covalent bond with a hydrogen ion, forming an ammonium ion. To make it clear that both electrons have been donated by ammonia, we use the same symbol for both electrons in the bond.

 
 

Average bond enthalpies are used to measure the strength of a covalent bond. Average bond enthalpy is the average amount of energy required to break the bond. Bonds with higher bond enthalpies are stronger than bonds with lower bond enthalpies since more heat energy is needed to break the bond. Bond enthalpies are approximations – the actual energy required to break a particular bond will depend on its surroundings (i.e. the other atoms and bonds present in the compound).

Shapes of molecules

The shapes of molecules depend on how many bonds they form and how many lone pairs they have. Bonding electrons and lone pair electrons will arrange themselves so that they are a maximum distance apart in order to minimise repulsion between them.

Lone pairs repel more than bonding pairs, so lone pair-lone pair bond angles are the largest, followed by lone pair-bond pair angles then bond pair-bond pair angles. This is because lone pairs are in orbitals which are shorter and rounder than the orbitals that bonding pairs occupy, so they take up more space.

If you know the number of bond pairs and lone pairs, then you can work out the shape of the molecule. For example, carbon dioxide has two bond pairs and no lone pairs so it must be a linear molecule with 180o bond angles. When considering the shapes of molecules you can treat double bonds as if they are single bonds. Look at the table below for the shapes associated with various combinations of bonding pairs and lone pairs.

 
 

Bond polarity and electronegativity

Electronegativity is a measure of the ability of an atom to pull an electron pair in a covalent bond towards itself. Fluorine is the most electronegative element, which means that out of all the elements in the periodic table, it is the best at attracting the electron pair in a covalent bond. Electronegativity can be measured on the Pauling scale where each element is assigned a Pauling value. The higher the Pauling value, the more electronegative the element.

For two atoms forming a covalent bond, the greater the difference in their electronegativities, the more polar the bond.

Although a molecule might have polar bonds, it can still be a non-polar molecule if it has a symmetrical geometry. When a molecule is arranged symmetrically, the polar bonds will cancel each other out.

Intermolecular forces

Intermolecular forces are the forces of attraction between different molecules. When a simple covalent substance, such as water, changes state it is the weak intermolecular forces that need to break, not the covalent (intramolecular) bonds within the molecule.

We can group intermolecular forces into two main categories:

  1. Permanent dipole-dipole interactions

  2. Induced (temporary) dipole-dipole interactions (aka London dispersion forces)

Permanent dipole-dipole interactions refers to situations where there is always a slight charge difference in the bond (polarity) resulting from the presence of an electronegative atom. The slightly positive region on one atom will be attracted to the slightly negative region on another atom of a different molecule.

 
 

Hydrogen bonding is a special type of permanent dipole-dipole interaction. It is the attraction between hydrogen on one molecule with an electronegative atom (usually oxygen, nitrogen or fluorine) on another molecule. Anything with an –OH (such as water and alcohols) can form hydrogen bonds. Hydrogen bonding is the strongest type of intermolecular force, which is why water has such a high melting and boiling point compared to similar liquids. Hydrogen bonding is also responsible for the fact that ice is less dense than liquid water. In the solid state there are more hydrogen bonds between water molecules. Since these bonds are fairly long, the water molecules are held a further distance apart compared to when they’re in the liquid state, making solid water less dense than liquid water.

 
 

Induced dipole-dipole interactions occur due to the random and uneven distribution of electrons within a molecule. Imagine a molecule of ethane with its electrons spinning around each of its atoms. If we froze it in time, there’ll be plenty of regions of the molecule where electrons happen to be present at that point in time. Likewise, there will be other regions that have no electrons. This uneven distribution of electrons temporarily creates regions that are slightly positive and other regions which are slightly negative. This is called an instantaneous dipole. When this molecule of ethane gets close to another ethane molecule, the electron-rich δ- regions will repel the electrons in the other molecule. This temporarily creates a dipole in that second molecule, which we call an induced dipole. The slightly negative region on one ethane molecule will be attracted to the slight positive region on the other. Another name given to induced dipole-dipole interactions is London dispersion forces.

Since electrons are responsible for induced dipole-dipole interactions, this means that ALL molecules experience these types of forces. The more electrons a molecule has, the stronger these interactions are. For example, as you go down group 7 boiling point increases because the halogen molecules have more electrons and therefore stronger London dispersion forces. This is why the halogens gradually change state from gaseous (F, Cl) to liquid (Br) to solid (I, At) as you descend the group.

Side note: you may have heard the term ‘van der Waals forces’. These refer to both permanent and induced dipole interactions. In the exam, I would avoid the use of this term to make it clear to the examiner that you know which type of intermolecular force you’re talking about.