Elements, Compounds and Mixtures

History of the Periodic Table

Mendeleev arranged his table in order of atomic mass, as well as using the properties of known elements and compounds.

Trying to order them based on mass is difficult, and sometimes led to untrue positions due to the fact that Mendeleev was unaware at the time about isotopes.

Where elements didn't fit the pattern, Mendeleev moved them so that his table made sense. This often led to there being gaps, and Mendeleev is most famous for leaving these gaps, claiming they were due to undiscovered elements. For some of these elements, he predicted their properties and masses... and got them right!

Mendeleev's Table (Key)

The symbols R²O and RH⁴, use superscripts to show the number of atoms in molecules rather than the current style of using subscripts. For group 1 ("Gruppe I") this would mean they could form molecules like H²O, and Li²O ... which we know today is correct. 

The gaps marked with hyphens show elements deduced by Mendeleev as existing, but unknown in 1872; he predicted the properties of some of these elements.

Mendeleev's Periodic Table

Modern Periodic Table

On the Periodic Table you will find every element that we know about, which in turn is a list of every atom that we have discovered. It is ordered by atomic number, meaning that reading from left to right increases the proton number by one every time.

The Periodic Table is arranged in such a way that elements with similar properties can be found together in columns called 'groups' - for example the group 1 metals are all very reactive with water.

Elements in the same group have the same number of electrons in their outershell (these are known as valence electrons).

Elements in the same period have the same number of electron shells.

The Periodic Table is split with metals all found on the left side of the table, and non metals found on the right.​ Metals tend to have 1-3 electrons on their outermost shell, so want to lose electrons to form positive ions (cations). Non-metals generally have 5-7 electrons on their outermost shell, so want to gain electrons to form negative ions (anions)

The Modern Periodic Table

Electron Configurations

Electrons orbit around an atom in a certain pattern. Around the nucleus are different energy levels, often called shells. Each shell can be filled to a specific number of electrons, and these limits are the same for every element. Each shell must be filled before another shell can take electrons.

Atoms will only fill to a maximum of:

  • 2 electrons in the first shell,
  • 8 in the second shell,
  • 8 in the third shell

We can either write, or draw, the electron configuration of an atom. Here you can see the electron configuration of potassium.


  • is in period 3 - so has 3 shells of electrons
  • is in group 1 - so has 1 electron on its outermost shell
  • has a configuration of because all shells must be filled before a new shell can contain electrons

Electronic Configuration of Potassium

Forming Ions

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

Common ions and their formulas

ion name example compound
oxide magnesium oxide (MgO)
hydroxide lithium hydroxide (LiOH)
F⁻, Cl⁻, Br⁻, I⁻
halide sodium chloride (NaCl)
nitrate potassium nitrate (KNO3)
carbonate calcium carbonate (CaCO3)
sulfate beryllium sulfate (BeSO4)

Forming Ions


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 outer shells 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.

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.

The Periodic Table is split with metals all found on the left side of the table, and non metals found on the right.​ Metals tend to have 1-3 electrons on their outermost shell, so want to lose electrons to form positive ions (cations). Non-metals generally have 5-7 electrons on their outermost shell, so want to gain electrons to form negative ions (anions).

Structure and Bonding in Metals

Non metals:

  • are dull (not shiny)
  • have low melting and boiling points
  • are poor conductors (good insulators)

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.

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 (dot and cross)

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.

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

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)

Carbon as an example

Carbon can form four covalent bonds, as it has four electrons on its outermost shell. The vast array of natural and synthetic organic compounds occur due to the ability of carbon to form families of similar compounds, chains and rings.

Covalent Bonding (dot and cross)

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.

Intermolecular forces between simple compounds

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.


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.

Structures of Diamond and Graphite

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 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.

Structures of Graphite, Graphene, Fullerens and Nanotubes.


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).


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 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

Bonding Models