The Periodic Table is an arrangement of elements in order of their atomic number. It shews clearly which elements have related properties. Elements are treated as members of families instead of as individuals, however, it must always be remembered that no one element is exactly like any other.
Elements within a Group or family shew:
The Periodic Table does not classify elements as metals or non-metals, however, there is a fairly obvious division between the two. Some elements are definitely metals whereas others are definitely non-metals. Those which are difficult to classify one way or the other are often called metalloids.
The 20 or so non-metals are found in the top right-hand corner of the Periodic Table.
On passing from the extreme left of the Periodic Table to the right there is a change from metal to metalloid to non-metal. On passing down a group in the Periodic Table there is an increase in metallic character of the element.
On going from left to right across the Periodic Table for the second and third periods, the melting point, boiling point, latent heat of fusion, latent heat of vaporisation and density shew an increase up to to a maximum at Group 4 and then fall to low values. This periodicity in physical properties depends on the structure and bonding of the elements.
Elements on the left exist as giant structures, whereas those on the right exist as discrete small molecules. Carbon has a a giant structure and hence has a high melting point, etc., whereas nitrogen exists as N2 molecules and hence has a low melting point, etc.
Across a period there is a gradual decrease in atomic size. This is because electrons are being added to the same energy level at about the same distance from the nucleus and protons are being added to the nucleus. Therefore, the electrons are attracted and pulled towards the nucleus with an increasing positive charge and so the radius of the atom decreases.
Within a period the first ionisation energy tends to rise with atomic number. This is due to the fact that the nuclear charge is increasing across the period from one element to the next and the electrons are added successively to the same energy level.
Electronegativity, the power of an atom to attract electrons to itself, increases across a period. This is because across a period the nuclear charge increases by one unit and one electron is added to the outer energy level. Since the nuclear charge increases, the atom has an increasing electron-attracting power.
The reactivity of elements changes across a period. Elements on the left are metals and tend to lose electrons to form positive ions, i.e. they are reducing agents. On the right elements tend to gain electrons to form negative ions, i.e. they are oxidising agents. Elements in the middle form covalent bonds.
There is a variation in the formula of the chlorides and oxides across periods 2 and 3.
|Moles Cl atoms / mole||1||2||3||4||5||2||-|
|Moles O atoms / mole||0.5||1||1.5||2||2.5||3||3.5|
The alkali metals and the alkaline earth metals are members of Group 1 and Group 2 respectively.
|Element||Symbol||Electronic Config||Flame Colour|
|Element||Symbol||Electronic Configuration||Flame Colour|
|Calcium||Ca||Ar 4s2||Brick red|
The elements in Group 1 all contain a single electron in their outer s orbital. In all compounds, Group 1 metals shew only an oxidation state of +1 which corresponds to the loss of this outer electron.
E.g. Na Na+ + e-
The elements in Group 2 all contain two electrons in their outer s orbital. Hence the oxidation state +2 exists in all their compounds. This corresponds to the loss of the two outer electrons.
In each group, the ionisation energies decrease and the chemical reactivities increase down the group. The decrease in ionisation energies down the groups is due to the increasing atomic radius and increased shielding of inner electrons down the group.
The elements are too reactive to occur in the native state. They are found as compounds with non metals.
In order to extract the s-block elements from their naturally occurring compounds, it is necessary to reduce their positive ions:
M+ + e- M
M2+ + 2e- M
The only way this can be practically carried out is by electrolysis of the molten electrolyte, e.g. sodium is obtained from sodium chloride.
All the metals in Groups 1 and 2 are high in the Activity (Electrochemical) series, because the outer s electrons are easily lost. Hence these metals are good reducing agents.
The difference in reactivity can be seen in the reactions of these metals with cold water. Lithium reacts slowly with cold water to produce the alkali, lithium hydroxide and hydrogen.
2Li(s) + 2H2O(l) 2LiOH(aq) + H2(g)
Sodium reacts more vigorously and potassium catches fire.
2K(s) + 2H2O(l) 2KOH(aq) + H2(g)
In Group 2, beryllium does not react with water. Magnesium reacts very slowly with hot water but rapidly with steam.
Mg(s) + H2O(g) MgO(s) + H2(g)
Calcium, strontium and barium react increasingly more rapidly with cold water.
Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)
The alkali metals and barium are usually stored in oil to prevent reaction with water and air.
A variety of oxides are produced when these metals burn in oxygen. For Group 1 metals the common oxide produced is M2O but peroxides and superoxides can be formed with the more reactive metals.
4Li(s) + O2(g) 2Li2O(s)
4K(s) + O2(g) 2K2O(s) + also K2O2 and KO2
Group 2 metals form the oxide MO, strontium and barium form peroxides MO2.
2Mg(s) + O2(g) 2MgO(s)
The oxides and hydroxides form alkaline solutions in water. (Oxides of non-metals form acidic solutions). The base strength increases down the Group.
The elements of Group 1 and Group 2 react with chlorine on heating to produce chlorides.
2Na(s) + Cl2(g) 2NaCl(s)
Mg(s) + Cl2(g) MgCl2(s)
The stability of a salt is dependent on both the size and the charge of its ions. The greater the charge, the stronger the attraction between the ions will be, and the more stable will be the compound. The smaller the ions become, the closer they can approach each other in the solid crystal, the more stable the compound.
When carbonates are heated they decompose to form the oxide. Group 1 carbonates are more stable than those formed from Group 2 metals. Sodium carbonate and potassium carbonate do not decompose. The carbonates become more difficult to decompose as you go down the group.
Magnesium carbonate is quite easily decomposed by heat, but calcium carbonate needs much stronger heating.
MgCO3(s) MgO(s) + CO2(g)
We say that the thermal stability increases down the group.
The pattern is similar for the nitrates as the carbonates. Similarly, compounds formed from Group 1 metals are more stable than those formed of Group 2 metals. The stability increases down the groups.
Group 1 nitrates decompose on heating to form the nitrite (except lithium nitrate).
2NaNO3(s) 2NaNO2(s) + O2(g)
Group 2 nitrates and lithium nitrate decompose on heating to form the oxide.
2Mg(NO3)2(s) 2MgO(s) + 4NO2(g) + O2(g)
The solubilities of ionic compounds of Group 1 and Group 2 depend on two factors:
Group 2 compounds where the anion has a single negative charge, e.g. hydroxide OH-, the solubility increases down the group. Magnesium hydroxide is less soluble than calcium hydroxide.
For Group 2 compounds where the anion has a double charge, e.g. carbonate CO32-, the solubility decreases down the group. Magnesium carbonate is more soluble than calcium carbonate.
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