Types of Bonding

Metallic

Metallic bonding is a form of bonding that holds together metals. To represent a metallic bond in a chemical equation, simply write the element's symbol. For example, solid gold would be ${\text{Au}}_{\left(\text{s}\right)}^{}$.

Metallic bonding consists of a number of metal ions in a lattice structure, with delocalised electrons moving within the lattice to hold the structure together. The electrostatic forces of attraction hold the lattice together.

The strength of a metallic bond is proportional to the metal's charge. Metals with a higher charge have more delocalised electrons in the lattice and higher forces of attraction, increasing the strength of the bond and therefore melting point.

The strength of a metallic bond is inversely proportional to the atomic radius. Metals with higher atomic radii will be further from the electrons, decreasing the strength of the bond and therefore melting point.

Metals have the following properties.

• High melting point: The forces of electrostatic attraction are quite strong, meaning that metals have high melting points.
• High electrical conductivity: Metals have free-moving charged particles, electrons, which can move through the lattice and conduct electricity.
• High thermal conductivity: The particles are very closely packed meaning not much energy is lost. In addition delocalised electrons can carry kinetic energy through the lattice structure.
• Malleable: The delocalised electrons can move and keep the structure together even if it is deformed, meaning metals are malleable.
• Ductile: Metals can be stretched out into wires for a similar reason to their malleability.

Ionic

Ionic bonding is a form of bonding that involves the complete transfer of electrons from one compound to another, giving them opposite charges, causing them to attract.

Ionic bonding does not always have to be between two monatomic ions; one example of an ionic bond that is not monatomic is ammonium chloride (${\text{NH}}_4^{}\text{Cl}$).

Ionic compounds can be split up into two compounds, solids and liquids, with different chemical properties.

Ionic Solids

Ionic solids consist of a number of ions arranged in a lattice structure. Ionic solids have the following properties:

• High melting point: The forces of electrostatic attraction are quite strong, meaning that ionic solids have high melting points.
• No electrical conductivity: The ions cannot freely move in the fixed lattice structure, meaning that ionic solids cannot conduct electricity.
• Thermal conductivity: The ions can move slightly from their normal positions, allowing the compound to conduct kinetic energy.
• Non-malleable: If a force is applied the layers of ionic compound move, meaning that like charges align and repel, shattering the compound.
• Non-ductile: See above.

Ionic Liquids

Ionic liquids consist of a number of ions in a liquid state.

Ionic liquids have the following properties:

• Electrical conductivity: The ions are able to move freely and conduct electricity.

Covalent

In covalent bonding atoms share electron pairs with one another so they have full valence electron shells.

There are two main types of covalent bonding.

Covalent Molecular

Covalent Molecular bonding is the most common type. In covalent molecular bonding, a number of atoms are combined to form individual molecules. For example, ${\text{CO}}_2^{}$ is covalently molecularly bonded.

Covalent molecular structures have the following properties.

• Low melting point: The only thing preventing the molecules from being melted is the intermolecular VDW (Van der Waals) forces, whose strength is proportional to the number of electrons. These require very little energy to overcome.
• Next to no electrical conductivity: Covalent molecular structures are usually not charged.

Covalent Network

Covalent Network bonding involves a number of atoms all covalently bonded to form what could be thought of as one extremely large molecule. There are two noteworthy network covalent elements: Carbon and Silicon.

Diamond is an allotrope of Carbon where every atom bonds with four others in a tetrahedral arrangement. Graphite is another allotrope of Carbon where every atom bonds with three others in a hexagonal arrangement, with delocalised electrons holding layers of Carbon atoms together.

• High melting point: The covalent bonds are very strong and require a large amount of energy to break.
• No electrical conductivity except Graphite: No structures have free-moving charged particles except for Graphite.

Comparing Melting Points

To compare the melting points of a set of elements, you must first find the types of bonding.

There are a few general rules we can use.

1. Covalent network $>$ Metallic, Ionic $>$ Covalent molecular (Since we are only discussing elements here, the Ionic is somewhat redundant, but may be useful if you are asked to compare the melting points of compounds).
2. Covalent molecular strength is proportional to number of electrons i.e. covalent molecular substances can be ordered by number of electrons.
3. Metallic strength is proportional to the charge of the metal ion. For example, ${\text{Al}}^{3+}$ has a higher melting point than ${\text{Na}}^{+}$.
4. You can use facts about atomic radius to sort metals. Generally a metal with a higher atomic radius will have a lower melting point, as the increased distance between metal atoms and the delocalised electrons will reduce the strength of the bond.
5. You can use knowledge of the states of elements at room temperature to deduce their relative melting points. For example we know that Oxygen is gaseous at room temperature, meaning it will have a lower melting point than an element which is solid at room temperature.

These are merely guidelines, and can all be deduced from the explanations of the individual types of bonding above (except rule 5 which is self-evident). Make sure that you explain what you have done in your answer.

As an example, we will use this question.

Q. Order the following elements by melting point, increasing: $\text{Ag},\text{S},\text{Hg},\text{Si},\text{P}$.

To answer this question we will first find the types of bonding used by each. Silicon is covalent network so it will have the highest melting point. Phosphorus and Sulfur are both covalent molecular so they will have the lowest melting points. Mercury and Silver are both metals so they will have melting points somewhere in the middle.

Phosphorus takes the form of ${\text{P}}_4^{}$ whereas Sulfur is ${\text{S}}_8^{}$. Since the melting point of a covalent molecular substance is proportional to the number of electrons, Sulfur will have a higher melting point than Phosphorus.

Mercury and Silver both have the same charge, so we cannot sort them as we usually would. We can use either rules 4 or 5 above to sort them. We know that Mercury is liquid at room temperature, and Silver is not, so Silver must have the higher melting point.

There is another step. So far we have assumed that the metallic elements have higher melting points than covalent molecular substances unconditionally. However, Mercury is the only element of the ones given that is a liquid at room temperature; all the others are solid. Therefore Mercury must have the lowest melting point of all the substances.

The final ordering is as follows:

A. $\text{Hg},\text{P},\text{S},\text{Ag},\text{Si}$.