Elements are arranged in the periodic table in order of increasing atomic number.

The periodic table allows chemists to make accurate predictions of physical properties and chemical behaviour for any element, based on its position.

Properties which follow a regular pattern across or down a periodic table are said to be periodic properties.

Features of the table are:

→ groups; groups are vertical columns within the table which contain elements with similar chemical properties resulting from a common number of electrons in the outer shell.

→ periods; periods are rows of elements in the table arranged with increasing atomic number, demonstrating an increasing number of outer electrons and a move from metallic to non-metallic characteristics.

The first 20 elements in the periodic table are categorised according to bonding and structure.

• The elements $Li$, $Be$, $Na$, $Mg$, $Al$, $K$, $Ca$ are metallic elements

  • Metal atoms are held together by metallic bonds and are arranged in a giant metallic lattice structure.

  

  • A metallic bond is the electrostatic force of attraction between the negatively charged, delocalised valence electrons and the positively charged nuclei of neighbouring atoms.
  • Metals are good conductors of both electrical and thermal energy as they have delocalised, charged electrons that can move freely through the lattice structure to transfer energy.
  • Metals are malleable and ductile because the electrons in a metallic lattice are delocalised and not fixed to a particular nucleus. Therefore, the metallic bonds in the lattice are not fixed and act in all directions, allowing metal atoms to slip past each other, relatively easily, to change the shape of the material.
  • Generally, metals have high melting and boiling points because the metallic bonds within the lattice structure that must be broken to boil or melt a metal are very strong, meaning a lot of energy is needed to overcome these bonds.
  • As you go across the periodic table, metals have more valence electrons that are delocalised in the metallic lattice. This means the electrostatic force of attraction is greater between them and the positively charged nucelli in the metal, as the ratio of delocalised electrons to positive nucleus has increased. This increases the strength of metallic bonds, so melting and boiling points are higher for these metals. These metals are also less malleable and ductile, as well as less reactive.
  • As you go down a group of metals, the atomic radius increases, meaning the delocalised valence electrons are farther away from their respective positive ions, and the screening effect due to inner shells of electrons is also greater. This means the electrostatic force of attraction between the delocalised electrons and positive ions is weaker. This makes the metallic bonds weaker, which means the melting and boiling points for these metals are lower. They are also more malleable and ductile, as well as very reactive.

• The elements $H_2$, $N_2$, $O_2$, $F_2$, $Cl_2$, $P_4$, $S_8$ and fullerenes — e.g. $C_{60}$ — are covalent molecular elements.

  • In general with the molecular elements, there are strong covalent bonds within discrete molecules — intramolecular forces — and weak London dispersion forces between molecules — intermolecular forces. The larger the molecule — i.e. the more atoms present in a molecule — the more electrons are present, the stronger the London dispersion forces $\therefore$ the higher the melting and boiling points of the element, as more energy is needed to break intermolecular bonds.

  • $H_2$, $N_2$, $O_2$ and the halogens exist as individual diatomic molecules and have weak London dispersion forces between the molecules — intermolecular forces — and strong covalent bonds between the atoms in the molecules — intramolecular forces. Due to only weak London dispersion forces between the molecules in these elements, they exist as gases at room temperature.

    • The diatomic molecules are all linear in shape.

    

    

    

  • The London dispersion forces in the larger sulfur and phosphorus molecules are stronger than in diatomic molecules as there are more electrons present in the former than the latter — as a greater number of atoms means more electrons in total in the molecule — so a greater temporary dipole can be created leading to a stronger electrostatic force of attraction between the positive and negative regions of molecules.

  • Sulfur and phosphorus are discrete molecules with stronger weak London dispersion forces between the molecules — intermolecular forces — and strong covalent bonds between the atoms in the molecules — intramolecular forces. As the London dispersion forces between sulfur and phosphorus molecules are stronger, they have higher melting and boiling points than the halogens and other elements that exist as diatomic molecules, and they exist as solids at room temperature.

    • Sulfur exists in molecules with 8 atoms — $S_8$. This is often referred to as a puckered ring or sulfur crown due to its shape.

      

    • Phosphorus exists in molecules with 4 atoms — $P_4$ — that have a pyramidal shape.

      

  • Carbon can exist as large discrete molecules known as fullerenes, with strong covalent bonds between the atoms within the molecule and weak London dispersion forces between the molecules. Fullerene molecules can be spherical (as $C_{60}$) or tube/plane shaped as graphene; they have relatively high boiling and melting points due to the stronger weak London dispersion forces between molecules as the molecules are so large.

    • $C_{60}$ is an example of a spherical fullerene, it made up of carbon atoms arranged in hexagons and pentagons.

    • Graphene is an example of a sheet shaped fullerene that can be rolled to make a tube, it is made up of carbon atoms arranged in hexagons.

      • In both of these structures each carbon atom has 3 valence electrons shared with another in covalent bonds and 1 free valence electron, however these electrons are not free to move between fullerene molecules, so, despite them, carbon in the form of these fullerenes does not conduct.

      

      Two on the left and bottom right are relevant to above information about carbon in the form of covalent molecules, top right is graphite which is in the next category of covalent network.

• The elements $B$, $Si$ and $C$ in the form of diamond and graphite are covalent networks.

  • Carbon can exist as two large covalent networks, graphite and diamond.

    • Carbon exists as a covalent network in the form of diamond.

      • In diamond, each carbon atom forms four strong covalent bonds with four other carbons in a strong three dimensional network of covalent bonds that makes diamond very strong and rigid.
      • The strong covalent network structure of diamond results in it having a very high melting and boiling point, as lots of strong covalent bonds must be broken to melt or boil it, which requires a lot of energy. Similarly, it is very difficult to cut or break diamond as it requires breaking many strong covalent bonds which requires a lot of energy.
      • Strong covalent bonds are formed in all three dimensions which means the bonds around each carbon atom are arranged in a tetrahedral shape.
      • All four valence electrons of each carbon are fixed in a covalent bond in diamond, leaving no free, delocalised electrons. This means diamond does not conduct electricity.

      

      Carbon in the form of diamond

    • Carbon also exists as a covalent network in the form of graphite.

      • In graphite, each carbon atom forms three strong covalent bonds with three other carbons in two dimensions, meaning the atoms are arranged hexagonally in planes.
      • Only three of each carbon’s valence electrons are used in covalent bonds, leaving a free valence electron for every carbon atom which is delocalised. Due to this presence of delocalised electrons that can move from atom to atom and carry a charge, graphite can conduct electricity.
      • The carbon atoms in graphite form hexagonal planes which are held together by weak London dispersion forces.
      • The layered arrangement of atoms in graphite makes it very brittle.
      • The layers in graphite can be easily separated, as only weak London dispersion forces have to be broken to accomplish this which requires less energy, therefore graphite is not as hard as diamond — it is easy to cut and break.

      

      Carbon in the form of graphite

  • Boron and silicon exist as covalent networks which means all atoms are held together by strong covalent bonds, resulting in high boiling and melting points.

    

    covalent network of boron atoms

    • Silicon is a tetrahedral covalent network.

      

    • Silicon has a slightly lower melting and boiling point than the other two elemental covalent networks as it has a greater number of occupied electron shells, meaning the valence electrons present in the covalent bonds of silicon are one shell farther away from the positive nucleus and the screening effect due to inner shell electrons is slightly greater. This means the electrostatic force of attraction between the negative bonding electrons and the positive nuclei in the covalent bonds is lower, making the covalent bonds slightly weaker.

• The noble gases — group 8 — are monoatomic elements.

  • The noble gases exist as single atoms with weak London dispersion forces between the atoms, meaning they have very low boiling and melting points and are all gases at room temperature.

The covalent radius is a measure of the size of an atom.

The covalent radius of elements increases down a periodic group as the number of occupied electron shells increases, which increases the overall size of the atom as the outer electrons are farther away from the nucleus and screened from its positive charge by an increased number of inner electron energy levels.

Although the number of protons also increases down a group, which generally decreases the covalent radius as there is an increased nuclear charge which pulls the electron shells closer which decreases the atom’s size, the increase in occupied electron shells, due to the screening effect of the increased number of inner electron shells, overpowers this.

The covalent radius decreases across a period as the increased number of protons in the nucleus of atoms leads to a greater nuclear charge meaning the occupied electron shells have a greater attraction to the nucleus and therefore held closer which decreases the atomic size.

There is no impact of number of occupied electron shells as this remains constant across a period, as the only increase in electrons across a period is within the same shell, which also means the screening effect of increasing the number of inner electron shells plays no part in diminishing the increased nuclear charge.

The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. The second and subsequent ionisation energies refer to the energies required to remove further moles of electrons.

→ i.e. the second ionisation energy is the energy required to remove two moles of electrons from one mole of gaseous atoms, the third ionisation energy is the energy required to remove three moles of electrons from one mole of gaseous atoms, etc.

→ NB: an element can only have as many ionisation energies as it does electrons

Ionisation energy equations can be written to show the the first ionisation energy and second ionisation energy of any element, etc., as follows.

→ NB: These must show one mole of the element in the gaseous form, as per the definition of the ionisation energy.

→ These equations can also be found printed on the top of page 12 of the data booklet.

$$ E_{(g)} \rarr E^+_{(g)}+e^- $$

$$ E^+{(g)} \rarr E^{2+}{(g)}+e^- $$

The first ionisation energy increases across a period, due to the increased positive nuclear charge which leads to a greater electrostatic attraction between the negative outer electron and the positive nucleus so more energy must be supplied to remove it.

The first ionisation energy decreases down a group, due to the increase in atomic size — as the number of occupied electron shells increases — meaning the outer electron is further away from the positive nucleus so the attraction between them is decreased and less energy is required to remove it. The increase in inner electron shells also leads to a screening effect which shields the negative outer electron from the positive nucleus and decreases the electrostatic attraction between them so less energy is needed to remove the electron.

Subsequent ionisation energies are generally higher than the first ionisation energy as the positive nuclear charge is shared between fewer electrons, so the remaining electrons experience a greater electrostatic attraction to the nucleus and more energy is required to remove them.