The Periodic Table

How the Periodic Table is arranged: groups, periods and proton number

The Periodic Table is an arrangement of all the elements in order of increasing proton number, which is also called the atomic number. Reading across a row the proton number goes up by one each time, so the table is really a long list of the elements written out in a grid.

The vertical columns are called groups and the horizontal rows are called periods. Elements in the same group have the same number of electrons in their outer shell, which is why they behave alike, and elements in the same period have the same number of occupied shells.

For the main groups (Groups I to VII) the group number is the number of outer shell electrons, and the period number is the number of occupied electron shells. Sodium has proton number 11 and configuration 2,8,1: three shells, so it is in Period 3, and one outer electron, so it is in Group I. Learn those two links and you can place any element up to proton number 20.

There is a second pattern running across a period. The elements on the left are metals and those on the right are non-metals, so there is a change from metallic to non-metallic character as you move across a period. The metals lose electrons to form positive ions, the non-metals gain or share electrons, and the elements in between have properties in between.

• Group = vertical column; for Groups I to VII the group number gives the number of outer shell electrons.
• Period = horizontal row; the period number gives the number of occupied electron shells.
• Elements are arranged in order of increasing proton number (atomic number).
• Metallic character decreases across a period, from metals on the left to non-metals on the right.

Group number, ion charge and predicting properties

The group number also tells you the charge on the ion an element forms. Atoms react so that they end up with a full outer shell. Group I atoms have one outer electron, so they lose it and form 1+ ions. Group II atoms lose two and form 2+ ions. Group VII atoms have seven outer electrons, so they gain one and form 1- ions, and Group VI atoms gain two to form 2- ions. Group VIII atoms already have a full shell, so they form no ions at all.

This is why the elements in a group have similar chemical properties. Their atoms all have the same number of outer shell electrons, so they lose, gain or share electrons in the same way. Every Group I metal reacts with water to give a hydroxide and hydrogen, and every Group VII halogen forms similar compounds with hydrogen and with metals.

Because of these patterns you can predict the properties of an element from its position. Given an unfamiliar element you can work out its electronic configuration from its group and period, predict the charge on its ion, and predict whether it will be a metal or a non-metal. The trends also repeat down each group, so a question that gives you information about two elements in a group is really asking you to spot the pattern and continue it.

The safest way to predict for an unfamiliar element is to compare it with the element above or below it in the same group. If astatine sits below iodine in Group VII, it will be a dark solid that is even less reactive than iodine, because that is the direction the Group VII trends run.

Group I: the alkali metals

The Group I elements lithium, sodium and potassium are called the alkali metals. They are relatively soft metals that can be cut with a knife, they have low densities so they float on water, and they are shiny when freshly cut but tarnish within seconds in air. They are stored under oil to keep the air and its water vapour away from the surface.

Three trends run down the group. The melting point decreases, so potassium is close to melting in hot water. The density increases, although all three metals are still less dense than water. The reactivity increases, so potassium reacts faster and more violently than sodium, and sodium faster than lithium.

The reaction with water is the classic test for this group. Lithium fizzes steadily and floats; sodium fizzes vigorously, melts into a silver ball and darts across the surface; potassium reacts so violently that the hydrogen it produces catches fire and burns with a lilac flame. In every case the general equation is: metal + water -> metal hydroxide + hydrogen.

For sodium the balanced equation is 2Na + 2H2O -> 2NaOH + H2. The solution left behind is alkaline, which is where the name alkali metal comes from. The reactivity increases down the group because the single outer electron is further from the nucleus and is shielded by more inner shells, so it is held less strongly and is lost more easily.

Group VII: the halogens

The Group VII elements are the halogens: fluorine, chlorine, bromine and iodine. They are diatomic non-metals, which means each molecule contains two atoms joined by a covalent bond, so their formulae are Cl2, Br2 and I2. They are the most reactive group of non-metals.

Their appearance changes strikingly down the group. At room temperature and pressure, chlorine is a pale yellow-green gas, bromine is a red-brown liquid and iodine is a grey-black solid. The colour deepens and the state changes from gas to liquid to solid as you go down, because the melting point and boiling point both rise.

Two trends matter. The density increases down the group. The reactivity decreases down the group, so chlorine is more reactive than bromine and bromine more reactive than iodine. This is the opposite direction to Group I, because a halogen atom needs to gain an electron to fill its outer shell, and the larger the atom the less strongly its nucleus attracts an incoming electron.

Chlorine and bromine have sharp, choking smells, and chlorine is toxic, so both must be handled in a fume cupboard. The name halogen means salt former, because these elements react with metals to form salts such as sodium chloride.

Displacement reactions of the halogens

A displacement reaction happens when a more reactive halogen takes the place of a less reactive halogen in a compound. If you bubble chlorine through a solution of potassium bromide, chlorine is more reactive than bromine, so it pushes the bromine out: chlorine + potassium bromide -> potassium chloride + bromine.

The balanced equation is Cl2 + 2KBr -> 2KCl + Br2. Written as an ionic equation it is Cl2 + 2Br- -> 2Cl- + Br2, which shows that the chlorine molecules gain electrons while the bromide ions lose them. The chlorine has been reduced and the bromide ions have been oxidised.

You can see the reaction happen because the colour changes. Potassium bromide solution is colourless, but after the reaction it turns orange or red-brown because bromine has been formed. In the same way chlorine displaces iodine from potassium iodide, turning the solution brown, and bromine displaces iodine from potassium iodide.

The rule is simple and worth memorising: a halogen can only displace a halide that is below it in the group, because only a more reactive halogen can take the place of a less reactive one. Iodine cannot displace bromine or chlorine, so adding iodine to potassium bromide solution gives no reaction and the solution stays colourless. These experiments are used to work out the order of reactivity of the halogens.

• Chlorine displaces both bromine and iodine.
• Bromine displaces iodine only.
• Iodine displaces neither chlorine nor bromine.
• No colour change means the halogen added is less reactive than the halide already in the solution.

The transition elements

The transition elements are the block of metals in the middle of the Periodic Table, between Groups II and III. They are typical metals, but their properties are quite different from the Group I metals. They have high densities and high melting points, so they are strong and hard-wearing, which is why they are used for structures, tools and machines.

Their compounds are coloured. Iron(II) compounds are usually pale green, iron(III) compounds are yellow or brown and copper(II) compounds are blue. This makes a useful test: a compound of a transition element is almost always coloured, whereas most compounds of Group I metals are white or colourless.

Transition elements and their compounds often act as catalysts. Iron catalyses the Haber process for making ammonia, vanadium(V) oxide catalyses the Contact process for making sulfuric acid, and manganese(IV) oxide catalyses the decomposition of hydrogen peroxide into water and oxygen.

The other key property is that transition elements have ions with variable oxidation numbers. Iron forms both iron(II), Fe2+, and iron(III), Fe3+, so the same element appears in two different compounds with different formulae. This ability to change oxidation number and change it back is part of the reason a transition element can act as a catalyst.

The noble gases

The noble gases are the elements of Group VIII, also called Group 0, at the far right of the Periodic Table: helium, neon, argon, krypton and xenon. They are colourless gases that are chemically unreactive, and they are monatomic, which means each particle is a single atom rather than a molecule.

The reason lies in their electronic configuration. Every noble gas has a full outer shell of electrons: helium has 2 outer electrons, neon has 2,8 and argon has 2,8,8. Atoms react in order to achieve a full outer shell, but a noble gas already has one, so it has no need to gain, lose or share electrons. With no bond to form, it stays as a single atom.

Because they do not react, the noble gases are useful whenever you need a gas that will not interfere with anything. Argon fills filament lamps and is used as a shield gas in welding because it will not react with the hot metal, helium fills balloons and airships because it is much less dense than air, and neon glows red-orange in advertising signs when a current is passed through it.

The reactivity of the noble gases does increase very slightly down the group as the atoms get larger, but for IGCSE purposes treat them all as unreactive. They form no ions and no compounds in the chemistry you are expected to know.

• Group VIII (Group 0): unreactive, monatomic, colourless gases.
• They have a full outer shell, so they do not need to gain, lose or share electrons.
• Helium has 2 outer electrons, neon 2,8 and argon 2,8,8.
• Uses: helium in balloons, argon in lamps and welding, neon in advertising signs.