Atomic Structure, Bonding and Periodicity - The Periodic Table
The Periodic table is an arrangement of elements in order of their atomic number, i.e. the number of protons in their nucleus.
Below is a selection of elements in the Periodic table, showing the different groups that have to be known at AS level,
Here is a link to a simple interactive full Periodic table:
There are also some very modern versions of the Periodic table, linking the elements more closely to the number and arrangement of electrons in their atoms. Here are two such links:
The modern version of the Periodic table classifies elements into groups (columns) and periods (rows).
Elements in the same group have a gradual change in physical properties such as melting point, boiling point and density as the group is descended and form compounds of similar chemical formulae.
Elements in the same period also show trends in physical properties, though they show much more variance in chemical reactivity, forming compounds with varying formulae along the period.
For example in the 3rd period, the electron configurations, electrical conductivities, atomic radii and melting/boiling points vary as follows -
(i) Variation in electron configuration/formulae -
|Atom||Electron configuration||Formula of oxide|
(ii) Variation of electrical conductivity -
The electrical conductivity increases from sodium to aluminium as these are metallic elements, with a sea of electrons in between nuclei; however, as the non-metal character of the elements takes over, the rest of the period are non-conductors (or semi-conductors in the case of silicon).
The electrons become tied up in covalent bonds - P4 molecules for phosphorus, S8 molecules for sulphur and Cl2 molecules for chlorine - and are therefore not available to conduct electricity.
(iii) Variation of atomic radius -
The atomic radius varies along period three as shown in the graph below,
The atomic radius decreases along the period as each atom adds one more electron to the same outer shell and one more proton to the nucleus. This increase in the number of protons increases the attractive pull of the nucleus and so the outer shell is pulled closer to the nucleus.
(iv) Variation of melting and boiling points -
The melting and boiling points vary along period three as shown in the graph below,
The atomi radii of the elements increase as the group is decended as the outer shell electrons are placed in a new shell. This shell is further away from the nucleus. Also, the (full) inner shells of electrons provide a shield from the attractive pull of the protons in the nucleus.
(i) Group II metals with oxygen -
|2M(s) + O2(g) 2MO(s)|
(ii) Group II metals with water -
|M(s) + H2O(l) M(OH)2(s) + H2(g)|
The reactivity of the metal with water increases as the group is descended, because the first and second ionisation energies decrease, meaning less energy is need to remove these two electrons and form a metal ion.
(iii) Magnesium compounds with hydrochloric acid -
|Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)|
|MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l)|
|MgCO3(s) + 2HCl(aq) MgCl2(aq) + H2O(l) + CO2(g)|
N.B.: The first reaction in this series is a redox reaction, i.e. the magnesium is oxidised and the hydrogen ions in the acid are reduced. The other two equations are acid-base reaction only - not redox.
(iv) Thermal decomposition of compounds -
|Oxides||Formula||DHf (kJ mol-1)|
|Carbonates||Formula||Td (oC)||DHf (kJ mol-1)||DHd (kJ mol-1)|
|Hydroxides||Formula||Td (oC)||DHf (kJ mol-1)||DHd (kJ mol-1)|
|Nitrates||Formula||Td (oC)||DHf (kJ mol-1)||DHd (kJ mol-1)|
(i) Limestone -
Calcium is a very common element in nature. It occurs most commonly in limestone rock as calcium carbonate, CaCO3. Limestone itself is used in many industrial processes, e.g. to make iron in the Blast Furnace, as a component of cement and in the manufacture of glass.
Limestone is commonly mines in open-cast mines, where the rock is blasted out of the ground, leaving massive pits. This is a relatively straight-forward and economically cheap way of mining; though it has a great impact on the local environment, with the pits requiring a lot of money to blend them into the local area once all the limestone has been extracted.
(ii) Quicklime -
When calcium carbonate is heated strongly in a slowly rotating kiln (see the diagram below)
it undergoes thermal decomposition to give calcium oxide, CaO, also known as quicklime, and carbon dioxide gas.
|calcium carbonate calcium oxide + carbon dioxide|
|CaCO3(s) CaO(s) + CO2(g)|
(iii) Slaked Lime -
A hydration reaction is one where water is added to a compound to make a new compound. When water is added to calcium oxide, calcium hydroxide, Ca(OH)2, is produced. This process used to be known as slaking lime and so calcium hydroxide is also known as slaked lime.
|calcium oxide + water calcium hydroxide|
|CaO(s) + H2O(l) Ca(OH)2(s)|
A dilute aqueous solution of calcium hydroxide is known as limewater. It has a pH of about 11. Limewater is used the laboratory to test for carbon dioxide gas. When carbon dioxide gas reacts with calcium hydroxide a white precipitate is formed. This precipitate is calcium carbonate
|calcium hydroxide + carbon dioxide calcium carbonate + water|
|Ca(OH)2(aq) + CO2(g) CaCO3(s) + H2O(l)|
(iv) Summary -
Calcium hydroxide in a solid form is used to neutralise acidity on farm land and in lakes affected by acid rain.
Magnesium hydroxide is used in indigestion/antacid tablets and solutions.back to top
The elements of Group 7 of the Periodic table are known as the halogens, from the classical Greek for salt forming. They are all non-metals, have different states at room temperature, different colours and a gradation of physical properties (see the table below).
|Element||State @ 25 oC||Colour||Melting pt. (oC)||Boiling pt. (oC)|
The halogens in general are toxic. Fluorine is particularly dangerous; you would already be dead by inhalation of fluorine before you could register its smell. Chlorine was used in the First World War as a poisonous gas, as it is heavier than air and could sink into trenches.
However, iodine is not toxic and can be used as a very effective disinfectant for wounds. Also, whilst solid iodine is black gaseous iodine has a distinctive purple colour.
The physical properties, i.e. melting and boiling points, of the halogens increase down the group, because as the group is descended, the diatomic molecules get bigger and there are more electrons in them. More electrons means greater Van der Waals forces and so a higher temperature is needed to overcome the forces.
The atomic radius increases as the group is descended as each new period adds an additional shell of electrons further away from the nucleus, which means the attractive pull of the nucleus is lessened. The inner shells also shield the outer shell from the attractive pull of the nucleus, weakening the force of attraction even more.
The reactivity of the halogens decreases as the group is descended. This is because the halogen atom needs to capture an electron to form an (negative) ion. With the smaller halogens, e.g. fluorine, the outer shell of electrons is close to the nucleus and so provides a greater attractive force to capture the electron. With the bigger halogens, e.g. iodine, the outer shell of electrons is much further away from the nucleus and so the electron is much harder to capture and ions form less readily.(2) Reactions :
(i) With metals -
All the halogens react with metals, though the ease of reaction decreases as the group is descended.
Just passing fluorine gas over reactive metals, such as the alkali metals, results in an exothermic reaction and the formation of a metal fluoride salt. However, with iodine the metal must be heated strongly before a reaction will occur.
Exemplar equation -
|aluminium + fluorine aluminium fluoride|
|2Al(s) + 3F2(g) 2AlF3(s)|
Try balancing this equation,
(ii) Production of chlorine -
When concentrated hydrochloric acid is added to solid potassium manganate(VII), KMnO4, chlorine gas is generated which can be collected using the apparatus shown.
N.B.: this must be done in a fume cupboard because of the toxic nature of chlorine.
The chemical equation for this reaction is not required at AS but is shown below,
|2KMnO4(s) + 16HCl(aq) 5Cl2(g) + 2MnCl2(aq) + 2KCl(aq) + 8H2O(l)|
The chlorine produced can be tested for by using damp litmus paper. Chlorine is a powerful bleaching agent and the litmus paper will be bleached, i.e. turn white. N.B.: this test will work for blue or red litmus paper, though blue litmus paper should turn red just before it is bleached, as an aqueous solution of chlorine is slightly acidic.
(iii) With other halides -
A halogen high in the group will displace a halogen lower in the group from that halogen's halide salt.
This can be demonstrated in the laboratory by adding some aqueous bromide and iodide ions to test tubes containing chlorine gas generated above.
When a solution of bromide ions are added the green gas turns orange-brown as does the solution, showing that elemental bromine is being formed.
Exemplar equations -
|magnesium bromide + chlorine magnesium chloride + bromine|
|MgBr2(aq) + Cl2(g) MgCl2(aq) + Br2(aq)|
With a solution of iodide ions, the green gas disappears and the solution turns a yellow colour with a black precipitate, i.e. insoluble solid, in it. The reason for these unusual observations is that iodine is slightly soluble in water and it forms a yellow solution when it does dissolve. Since most of the iodine formed does not dissolve the rest forms the black solid.
Exemplar equations -
|magnesium iodide + chlorine magnesium chloride + iodine|
|MgI2(aq) + Cl2(g) MgCl2(aq) + I2(s)|
These reactions can also be represented by what is known as an ionic equation. This type of equation does not show any elements that do not change during the reaction. So for the reactions above of chlorine with bromide and iodide ions, the ionic equations would look like this -
|2Br-(aq) + Cl2(g) 2Cl-(aq) + Br2(aq)|
|2I-(aq) + Cl2(g) 2Cl-(aq) + I2(s)|
As can be seen this representation can simplify dramatically a chemical equation by removing the metal ions.
Try balancing these equations using fluorine and some halide salts,
(iv) Reactions of halide ions -
Experimental sheet for the reactions of halide ions.
Halide ions undergo a series of unique reactions that allow an unknown solid or aqueous sample to be tested for the presence of chloride, bromide or iodide ions.
Aqueous silver ions react with halide ions to produce individually coloured precipitates. These precipitates have different solubilities in ammonia solution and so further differentiation can be achieved, see the table below,
|Halide ion||Observation with AgNO3(aq)||Solubility in NH3(aq)|
|Chloride||white precipitate||soluble in dilute NH3(aq)|
|Bromide||cream precipitate||soluble in conc. NH3(aq)|
|Iodide||pale yellow precipitate||not soluble in NH3(aq)|
Aqueous lead(II) ions react with halide ions to produce a different set of halide salts, see the table below,
|Halide ion||Observation with Pb(NO3)2(aq)|
|Iodide||bright yellow precipitate|
Concentrated acids react with solid halide salts to produce a variety of results, depending on the acid. Strong oxidising acids, such as concentrated sulphuric acid, form a variety of products depending on the halide used. With chloride ions, only hydrogen chloride is formed; with bromide ions, bromine, sulphur dioxide and water are formed; and with iodide ions, iodine, hydrogen sulphide and water are formed.
A normal strong acid, such as phosphoric(V) acid, only forms the hydrogen halide and a phosphate(V) salt.(3) Industrial Resources :
Sodium chloride is a great natural resource, present in great abundance in sea water and underground lakes, as a concentration aqueous solution called brine. This brine also contains a myriad of other compounds, including bromide and iodide salts as well as more gold than is present on land.
The sodium chloride can be obtained from sea water by simply evaporating the water from the solution and collecting the solid left behind. This generally requires a lot of energy to evaporate the water so is not always economically viable.
The sodium chloride solution can be used to make pure sodium metal or sodium hydroxide (a powerful caustic base) and chlorine gas itself.
Solid sodium chloride, along with other salts, is spread on icy roads, as it lowers the freezing (or melting) point of water, thus causing any ice already present to melt. This can only work at temperatures higher than -10 oC or so, otherwise other salt mixtures may be needed.
Chlorine by itself is used as a bleach and in the manufacture of sodium chlorate, which can be used as a bleach and a herbicide.
Water purification also relies on chlorine to kill bacteria in the water, after the impure water has passed through various filtration stages.
Chlorine is also used in the production of chlorofluorocarbons, commonly called CFC's, used in the past as refrigerant gases and propellants for aerosol cans. Both these uses have now been banned by international law, in the developed world at least.
The problem these chemicals cause is that when they reach the high atmosphere the molecules break apart to release chlorine atoms. These chlorine atoms then react with molecules of ozone, O3, turning them into molecules of oxygen. One chlorine atom can destroys thousands of molecules of ozone.
Ozone traps UV light from the sun, preventing it hitting the surface of the Earth, and with a depletion of ozone more UV light gets through, which increases occurrences of skin cancer in humans.
Unfortunately, in place of CFC's, alkanes are used and these lead to an increase in the amount of green house gases and an increase in global temperatures.back to top
written by Dr Richard Clarkson : © Saturday, 1 November 1997
Updated : Tuesday, 16th August, 2011
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