Periodic Table
The Periodic Table is a vital tool used by chemists to predict the way in which elements react during chemical reactions. It is a method of categorising elements according to their properties.
The period number tells you how many shells there are. All the elements in a group have the same number of electrons in their outer shells. So Group I elements have 1, Group II have 2, and so on. These outer-shell electrons are also called the valency electrons. The group number is the same as the number of outer-shell electrons. The valency electrons dictate how an element reacts. So the elements in Group I all have similar reactions, for example.
Example
The element calcium, Ca, is the 20th element in the Periodic Table. It is in Group II, Period 4. Its proton number is 20. What is its electron distribution?
Group II tells you there is 2 electrons in the outer shell. Period 4 tells you there are 4 shells. The proton number is 20, so there are also 20 electrons.
Elements with similar chemical properties are found in the same groups. There are eight groups of elements. The first column is called Group I; the second Group II; and so on up to Group VII. The final column in the Periodic Table is called Group 0 (or Group VIII).
Group names
- Group I: The alkali metals
- Group II: The alkaline earth metals
- Group VII: The halogens
- Group 0 (VIII): Inert gases or noble gases
| Group Number | ||||||||
|---|---|---|---|---|---|---|---|---|
| 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | |
| Period 1 | 11H | 11H | 24He | |||||
| Period 2 | 37Li | 49Be | 511B | 612C | 714N | 816O | 919F | 1020Ne |
| Period 3 | 1123Na | 1224Mg | 1327Al | 1428Si | 1531P | 1632S | 1735.5Cl | 1840Ar |
| Period 4 | 1939K | 2040Ca | ||||||
Group I – the alkali metals
They include Lithium, sodium and potassium. They are all very reactive metals and they are stored under oil to prevent them coming into contact with water or air.
These three metals have the following properties
- They are good conductors of electricity and heat.
- They are soft metals. Lithium is the hardest and potassium the softest.
- They are metals with low densities. For example, lithium has a density of 0.53 g cm–3 and potassium has a density of 0.86 g cm–3. So they float on water while reacting with it.
- They have shiny surfaces when freshly cut with a knife
- They have low melting points. For example, lithium has a melting point of 181 °C and potassium has a melting point of 64 °C.
- They burn in oxygen or air, with characteristic flame colours, to form white solid oxides. For example, lithium reacts with oxygen in air to form white lithium oxide, according to the following equation:
- These Group I oxides all dissolve in water to form alkaline solutions of the metal hydroxide e.g.
- They react vigorously with water to give an alkaline solution of the metal hydroxide as well as producing hydrogen gas e.g.
- They react vigorously with halogens, such as chlorine, to form metal halides, for example sodium chloride.
lithium + oxygen → lithium oxide
4Li(s) + O2(g) → 2Li2O(s)
lithium oxide + water → lithium hydroxide
Li2O(s) + H2O(l) → 2LiOH(aq)
potassium + water → potassium hydroxide + hydrogen gas
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
sodium + chlorine → sodium chloride
2Na(s) + Cl2 (g) → 2NaCl(s)
Of these three metals, potassium is the most reactive towards water, followed by sodium and then lithium. The further down the group you go the more reactive the metals become. Francium is, therefore, the most reactive Group I metal
| Element | Symbol | Proton number | Electronic structure |
|---|---|---|---|
| Lithium | Li | 3 | 2,1 |
| Sodium | Na | 11 | 2,8,1 |
| Potassium | K | 19 | 2,8,8,1 |
Sodium atom
Potassium atom
Group II – the alkaline earth metals
Magnesium and calcium are the most common group 2 metals.
These metals have the following properties.
- They are harder than those in Group I.
- They are silvery-grey in colour when pure and clean.
- They tarnish quickly, however, when left in air due to the formation of a metal oxide on their surfaces
- They are good conductors of heat and electricity.
- They burn in oxygen or air with characteristic flame colours to form solid white oxides. For example:
- They react with water, but they do so much less vigorously than the elements in Group I. For example:
magnesium + oxygen → magnesium oxide
2Mg(s) + O2(g) → 2MgO(s)
calcium + water → calcium hydroxide + hydrogen gas
Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
The further down the group you go, the more reactive the elements
become
| Element | Symbol | Proton number | Electronic structure |
|---|---|---|---|
| Beryllium | Be | 4 | 2,2 |
| Magnesium | Mg | 12 | 2,8,2 |
| Calcium | Ca | 20 | 2,8,8,2 |
Magnesium atom
Calcium atom
Group VII – the halogens
Group VII includes elements such as fluorine, chlorine, bromine and iodine.
These elements are coloured and darken going down the group.
| Halogen | Colour |
|---|---|
| Chlorine | Pale green Image source |
| Bromine | Red–brown Image source |
| Iodine | Purple–black vapour/ Silver-grey solid Image source |
Halogens have the following properties
- They exist as diatomic molecules, for example Cl2,
Br2 and I2. - They show a gradual change from a gas (Cl2), through a liquid (Br2), to a solid (I2) as the density increases.
- They form molecular compounds with other non metallic elements, for example HCl.
- They react with hydrogen to produce the hydrogen halides, which dissolve in water to form acidic solutions.
- hydrogen + chlorine → hydrogen chloride
- H2(g) + Cl2(g) → 2HCl(g)
- hydrogen chloride + water → hydrochloric acid
- HCl(g) + H2O → HCl(aq)
| Element | Symbol | Proton number | Electronic structure |
|---|---|---|---|
| Fluorine | F | 9 | 2,7 |
| Chlorine | Cl | 17 | 2,8,7 |
| Bromine | Br | 35 | 2,8,18,7 |
Uses of halogens
- Fluorine is used in the form of fluorides in drinking water and toothpaste to reduce tooth decay by hardening the enamel on teeth.
- Chlorine is used to make PVC plastic as well as household bleaches. It is also used to kill bacteria and viruses in drinking water
- Bromine is used to make disinfectants, medicines and fire retardants.
- Iodine is used in medicines and disinfectants and
also as a photographic chemical.
Group 0 – the noble gases
Also known as group 8, includes elements such as Helium, neon, argon, krypton, xenon.
Their atoms all have 8 outer-shell electrons, except for helium, which has 2 because it has only one shell. This stable arrangement of electrons has a very important result because it makes the Group 0 elements unreactive.
Noble gasses have the following properties
- They are colourless gases.
- They are monatomic gases – they exist as individual atoms, for example He, Ne and Ar.
- They are very unreactive.
| Element | Symbol | Proton number | Electronic structure |
|---|---|---|---|
| Helium | He | 2 | 2 |
| Neon | Ne | 10 | 2,8 |
| Argon | Ar | 18 | 2,8,8 |