Core Practicals

Structure and Bonding

Forming Ions

calcium carbonate
Cl⁻, Br⁻, I⁻
sodium chloride
lithium hydroxide
potassium nitrate
magnesium oxide
beryllium sulfate
Forming Ions (e.g. fluoride)

An ion is an atom (or group of atoms) with a positive or negative charge, formed by either losing or gaining electrons.


Electrons are negatively charged, and protons are positively charged. Atoms are neutral overall, due to the fact they have equal numbers of protons and electrons. By changing the number of electrons (either by an atom losing, or gaining electrons) - the atom forms an ion.


Metals generally form positive ions as it is easier for them to lose electrons to reveal a full shell of electrons, and non metals generally form negative ions as it is easier for them to gain electrons to make a full shell:

  • for example a group 2 metal will form a 2+ ion by losing two electrons, forming a positive ion (cation); we say that it has been oxidised

  • for example a group 6 non-metal will form a 2- ion by gaining two electrons, forming a negative ion (anion); we say that it has been reduced

Ionic Bonding

Ionic compounds can be described as having a lattice structure that consists of a regular arrangement of ions held together by strong electrostatic forces between oppositely charged ions. Ions form when atoms transfer electrons (one atom loses electrons, the other gains).


Ionic substances:

  • have high melting points, as large amounts of energy are needed to break the electrostatic forces of attraction

  • can conduct electricity when molten or dissolved in water, as the ions are free to move (the charge can flow)

  • cannot conduct electricity when solid, as there are no free moving ions


When drawing a dot-and-cross diagram for ionic bonds you must remember to include the following:

  1. square brackets to show that it ions are formed

  2. the charge of each ion to the top right of the square brackets

  3. only the outershell electrons (also called valence electrons) need to be drawn (unless otherwise stated)

Ionic bonding (NaCl)

Rules for naming compounds

If the compound has more than one part to its name (e.g. sodium chloride), then the element furthest to the left in the Periodic Table comes first in the name.

Elements Involved
a metal, a non-metal and oxygen
magnesium sulfate (MgSO₄)
a metal and one non-metal
magnesium sulfide (MgS)

Covalent Bonding

A covalent bond is formed when a pair of electrons is shared between two non-metal atoms. Each atom must share 1 electron each to form one covalent bond - so each covalent bond is made of 2 electrons.


When drawing a dot-and-cross diagram for covalent bonds you must remember:

  1. only the outershell electrons (also called valence electrons) need to be drawn (unless otherwise stated)

  2. each covalent bond is a pair of electrons (one electron from each atom in the bond)

  3. there can be double bonds (4 shared electrons), or even triple bonds (6 shared electrons)

Covalent Bonding

A substance that contains atoms held together by covalent bonds is referred to as a molecule. Simple molecules are typically around 0.1 nanometers in size.

Simple Molecular Compounds

Simple molecular compounds (simple molecules) contain only a few atoms, and we can tell how many of each atom is in the molecule by looking at its formula.


These compounds are usually liquids or gases at room temperature as the molecules are held together by weak intermolecular forces of attraction (but the atoms in the compounds are held together internally by strong covalent bonds), so only a small amount of energy is required to change state. This means simple molecules often have low melting and boiling points.


Simple molecules do not conduct electricity as there are no free moving electrons or ions.

Weak intermolecular forces (simple covalent bonds)

Diamond and Graphite

Graphite and diamond are examples of giant covalent structures. These compounds are solid at room temperature, because all of the atoms in a giant covalent structure are held together by strong covalent bonds. These bonds have to be broken by large amounts of energy leading to high melting and boiling points.

Diamond and Graphite

Graphite is made from layers of hexagonal rings of carbon, with each atom forming three strong covalent bonds to other carbon atoms. Each atom has a 'spare' electron, not used for bonding, which it contributes to the “sea of delocalised electrons”, thereby being able to conduct heat and electricity well.​ This is why graphite is often used for electrodes in electrolysis.


Weak forces of attraction hold the layers of graphite together, so they can slide over each other, making graphite a great lubricant.


Every carbon atom is strongly covalently bonded to four others in diamond, and because of this it forms a 3D lattice, called a tetrahedron. No free electrons exist in this structure, so it does not conduct electricity. 


Diamonds are used as cutting tools as they are the hardest naturally occurring substance due to the arrangement of carbon atoms bonded covalently.

Other Allotropes of Carbon

Carbon can form many different structures with different properties, and when elements can do this - we call them allotropes.


Graphene is just a one atom thick (single layer) of graphite. It also contains free moving electrons, and so is very good at conducting electricity.


Fullerenes (such as C₆₀) can be thought of as graphene sheets rolled into a ball, however graphene is made of 6 sided rings, and fullerenes are made of 5 and 6 sided rings. Fullerenes can take the shapes of balls, or other shapes like tubes (nanotubes). 


Fullerenes and nanotubes have similar properties to graphene, however nanotubes have a high tensile strength.

10 allotropes-01.png


Polymers are large molecules, made of ‘repeating units’ called monomers.


All the atoms in a polymer are bonded to other atoms to make a long chain of strong covalent bonds, usually with a carbon backbone.


Because polymers can be very long, we don't write out their full structure - and instead we can show their repeating units in a diagram similar to the one shown. This shows how we can turn ethene into poly(ethene).

11 polymers-01.png

Metallic Bonding

Metals have some unique properties - and it's all to do with how they bond! Their structure is formed from positive metal ions held together by a “sea of delocalised electrons” from the outershells of the metal atoms. The electrons are free to move around from each atom to atom as they please.


The strong electrostatic forces between the ions and electrons mean metals have very high melting points (large amounts of energy are needed to break these forces).


Because the electrons are able to move freely, it means metals are good conductors of electricity and heat.


Metals are also shiny, as well as malleable (bendable) and ductile (can be drawn into wires) as they have regular layers of atoms. These layers can slide over each other if they are hammered.

12 metallic bonding-01.png

Alloys are less malleable than pure metals as they have irregular layers, and so they cannot slide over each other as easily. The atoms are still held together by metallic bonding.


Non metals:

  • are dull (not shiny)

  • have low melting and boiling points

  • are poor conductors (good insulators)

Bonding Models

We use models to help us represent what molecules look like.


Ball and Stick

This model allows us to see the arrangement of atoms in 3D, but does not give us an accurate image of how much space the atoms take up (and bonds aren't really lines!).​


Space Filling (covalent) or Close Packed (ionic)

This model allows us to see the 'true space' that atoms (or ions) take up, and how much overlap of electron density there is. However, we cannot always rely on this model as it does not show charges, and sometimes it is hard to see which atoms are bonded.


Straight Lines (covalent)

This is the most common way to represent covalent bonds, with each line representing a shared pair of electrons (covalent bond). These diagrams allow us to see which atoms are bonded, but not the size of atoms, or their electron densities.


Examples include:

  • hydrogen gas              H-H

  • carbon dioxide         O=C=O

  • water                          H-O-H

13 shapes of molecules-01.png