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Groups and trends in Physical Properties
Transcript of Groups and trends in Physical Properties
The boiling points of these alkali metals show a similar pattern to the melting points. Group 2 Appearance
The Group 2 elements are all metals with a shiny, silvery white colour
The Alkaline Earth Metals are high in the reactivity series of metals, but not as high as the Alkali Metals of Group 1.
Occurrence and Extraction
These elements are all found in the Earth's crust, widely distributed in rock structures in their non-elemental forms. Only Magnesium is extracted on a large scale. It is extracted from sea water by the addition of Calcium hydroxide, which precipitates out the less soluble Magnesium hydroxide. This is then converted to Magnesium chloride, which is then electrolysed in a Downs Cell.
The metals of Group 2 are harder and denser than Sodium and Potassium, and have higher melting points. These properties are due to the presence of two valence electrons on each atom, which gives stronger metallic bonding than occurs in Group 1.
These elements give characteristic colours when heated in a flame:
Magnesium - Brilliant White
Calcium - Brick Red
Strontium - Crimson
Barium - Green
Radium - Red
Atomic and ionic radii increase down the Group, but ionic radii are smaller than the corresponding atomic radii. This is because the atoms lose their two outer electrons to form ions. The remaining electrons occupy closer levels, and the increased effective nuclear charge attracts the electrons towards the nucleus.
The chemical properties of Group2 elements are dominated by the strong reducing power of the metals. Once started, the reactions with Oxygen and Chlorine are vigorous:
2Mg(s) + O2(g) è2MgO(s)
Ca(s) + Cl2(g) è CaCl2(s)
All the metals except Beryllium form oxides in air at room temperature which dulls the surface of the metal. Beryllium has to be stored under oil.
All the metals, except Beryllium, reduce water and dilute acids to Hydrogen:
Mg(s) + 2H+(aq) è Mg(aq) + H2(g)
Magnesium reacts only slowly with water unless the water is boiling, but Clacium reacts rapidly even at room temperature, and forms a cloudy white suspension of sparingly soluble Calcium hydroxide.
Calcium, Strontium and Barium can reduce Hydrogen gas when heated, forming the hydride:
Ca(s) + H2(g) è CaH2(s)
The hot metals are also sufficiently strong reducing agents to reduce Nitrogen gas and form nitrides:
3Mg(s) + N2(g) è Mg3N2(s)
Magnesium can reduce, and burn in, Carbon dioxide:
2Mg(s) + CO2(g) è 2MgO(s) + C(s)
This means that Magnesium fires cannot be extinguished using Carbon dioxide fire extinguishers.
The oxides of Group 2 metals have the general formula MO and are basic. They are normally prepared by heating the hydroxide or carbonate. They have high melting points. Peroxides, MO2, are known for all these elements except Beryllium
Calcium, Strontium and Barium oxides react with water to form hydroxides;
CaO(s) + H2O(l) è Ca(OH)2(s)
Calcium hydroxide is known as Slaked Lime. It is sparingly soluble in water and the resulting mildly alkaline solution is known as Limewater which is used to test for the gas Carbon dioxide.
The Group 2 halides are normally found in the hydrated form. They are all ionic except for Beryllium chloride. Anhydrous Calcium chloride has such a strong affinity for water it is used as a drying agent. Melting point
The alkali metals have low melting and boiling points compared to most other metals. Apart from the other alkali metals, only three metals (indium, gallium and mercury) have lower melting points than lithium. You can see from the graph that lithium, at the top of Group 1, has the highest melting point in the group. The melting points then decrease as you go down the group The density of a substance is a measure of how much mass it has for its size. It is measured in grams/cubic centimetre. For example gold and lead are very dense metals - even a small lump of either of them can still feel heavy. The alkali metals have low densities compared to most other metals. (They feel lighter.) You can see from the graph that lithium, at the top of Group 1, has the lowest density in the group. The densities then generally increase as you go down the group. Group 3 Group 6 General Information Columns ON THE PERIODIC TABLE ARE CALLED GROUPS. Elements within a group share A FEW common properties. Elements IN A GROUP have the same outer electron arrangement. elements in a group share similar chemical properties. Group 7 Has halogens have low melting points and boiling points. This is a typical property of non-metals. Fluorine has the lowest melting point and boiling point. The melting points and boiling points then increase as you go down the group. As you go down a Group they will have a low ionization energy because the loss of an electron forms a stable shell . It becomes harder to remove an electron as the atomic radius decreases because the electrons are generally closer to the nucleus, which is also more positively charged. Few Compounds from Group 1
3) Sodium Hydrogen Carbonate - NaHCO3.
Called 'bicarb', is used in cooking as a raising agent for cakes.
Baking powder is a mixture of sodium hydrogen carbonate
and tartaric acid, to keep the pH neutral. Sodium Chloride - NaCl.
Sodium chloride is common salt.
It is used in the food industry as a flavouring and as a preservative.
2) Sodium Carbonate - Na2CO3.
Sodium carbonate (called soda, or washing soda)
is used in the manufacture of glass
and to make water soft, replacing calcium ions with sodium ions. 4) Sodium Hydroxide (caustic soda) - NaOH. The atomic radius increases as you go down the Group Has a decreasing atomic radius Increasing Ionisation energy. Physical Properties of Group 2 Increases down each group electrons are in shells further from the nucleus Ionic Size: Increases down the group. Melting Points: Decrease down each group metallic bonding gets weaker due to increased size. Ionisation Energy Decreases down the group atomic size increases. Li =
Cs= Red strong persistent orange Lilac (pink) Red (reddish-violet) Unknown Trends in Melting points and Boiling points in Period 3
Melting points increase from sodium to silicon, then decrease going to argon (with a “bump” at sulphur). Boiling points generally increase going from sodium to aluminum, then decrease to argon (again with a “bump” at sulphur). When a substance melts, some of the attractive forces holding the particles together are broken or loosened so that the particles can move freely around each other but are still close together. The stronger these forces are, the more energy is needed to overcome them and the higher the melting temperature.
When a substance boils, most of the remaining attractive forces are broken so the particles can move freely and far apart. The stronger the attractive forces are, the more energy is needed to overcome them and the higher the boiling temperature. The metals of Group 3 are harder and denser than Group 2. Which increases as you go down. , the have a have increasing melting and Boiling points. F http://uploadpic.org/storage/2011/a8H5DhO2IK9E1tLzsHQhLgUFl.gif
All the Group 1 elements are silvery coloured metals. They are soft, and can easily be cut with a knife to expose a shiny surface which dulls on oxidation.
These elements are highly reactive metals. The reactivity increases on descending the Group from Lithium to Caesium. There is a closer similarity between the elements of this Group than in any other Group of the Periodic Table.
Occurrence and Extraction
These elements are too reactive to be found free in nature. The alkali metals are so reactive they cannot be displaced by another element, so are isolated by electrolysis of their molten salts.
The alkali metals are soft, with low melting and boiling temperatures. The have low densities - Lithium, Sodium and Potassium are all less dense than water. They all show weak metallic bonding as only one electron is available from each atom.
Alkali metals colour flames.
Lithium - Red
Sodium - Yellow
Potassium - Lilac
Rubidium - Red
Caesium - Blue
The ionic radii of the alkali metals are all smaller than the atomic radii. This is because the atom contains one outer electron which is lost to form the ion. The remaining electrons are all in levels closer to the nucleus. In addition, the increased effective nuclear charge attracts the remaining electrons towards the nucleus and decreases the size of the ion.
The alkali metals are all strong reducing agents. They can reduce Oxygen, Chlorine, Ammonia and Hydrogen. The reaction with Oxygen tarnishes the metals in air, so they are all stored under oil. They cannot be stored under water because they react with it to produce Hydrogen and alkali hydroxides;
2Na(s) + 2 H2O(l) è 2NaOH(aq) + H2(g)
Lithium dissolves steadily in water with effervescence; Sodium reacts more violently and can burn with an orange flame; Potassium ignites on contact with water and burns with a lilac flame; Caesium sinks in water, but the rapid generation of Hydrogen gas produces a shock wave that can shatter a glass container.
Sodium dissolves in liquid Ammonia to give a deep blue solution. This solution is used as a reducing agent. At higher concentrations the colour of the solution changes to bronze and it conducts electricity like a metal.
The chemistry of Lithium shows some anomalies. As the Li+ ion is so small it polarises anions and so introduces a covalent character to its compounds. It has some relationship to Magnesium.
The alkali metals form ionic solid oxides of composition M2O when burnt in air. However, Sodium forms the peroxide Na2O2 as the main product, and Potassium forms the superoxide KO2, also as the main product.
Alkali metal hydroxides are white ionic crystalline solids of formula MOH, and are soluble in water. They are all deliquescent except LiOH. The aqueous solutions are all strongly alkaline. They neutralise acids to form salts;
NaOH(aq) + HCl(aq) èNaCl(aq) + H2O(l)
Alkali metal halides are white ionic crystalline solids. They are all soluble in water with the exception of LiF, re Appearance
Boron is a non-metallic grey powder, and all the other memebers of the Group are soft, silvery metals. Thallium develops a bluish tinge on oxidation.
The general trend down Group 3 is from non-metallic to metallic character. Boron is a non-metal with a covalent network structure. The other elements are much larger than Boron and are more ionic and metallic in character. Aluminium has a close-packed metallic structure but is on the borderline between ionic and covalent character in its compounds. The remainder of Group 3 are generally considered to be metals, although some compounds show covalent characteristics.
Occurrence and Extraction
These elements are not found free in nature, but are all present in various minerals or ores. Aluminium is the most widely used element in this Group. It is obtained by the electrolysis of Aluminium oxide, which is purified from Bauxite.
The influence of the non-metallic character in this Group is reflected by the softness of the metals. The melting points of all the elements is high, but the melting point of Boron is much higher than that of Beryllium in Group 2, whereas the melting point of Aluminium is similar to that of Magnesium in Group 2. The densities of all the elements in Group 3 are higher than those in Group 2.
The ionic radii of the elements are much smaller than the atomic radii, as three outer electrons are lost in the formation of the ions. The resulting increased effective nuclear charge attracts the remaining electrons closer to the nucleus.
The chemical properties of the elements of Group 3 reflect the increasingly metallic characteristics of descending members of the Group. Only Boron and Aluminium will be considered here.
Boron is unreactive except at high temperatures. Alunminium is a highly reactive metal which is readily oxidised in air. This oxide coating is resistant to acids but is moderately soluble in alkalis. Aluminium can reduce strong alkali, a product being the tetrahydroxyaluminate ion, Al(OH)4-. Aluminium also reacts violently with Iron (III) oxide to produce Iron in the Thermit process:
2Al(s) + Fe2O3(s) è 2Fe(s) + Al2O3(s)
Boron oxide is an insoluble white solid with a very high boiling point (over 2000K) because of its extended covalently bonded network structure. Aluminium oxide is amphoteric.
Boron forms an extensive series of Hydrides, the Boranes. The simplest of these is not BH3 as expected, but its dimer B2H6
The most important halide of Boron is Boron trifluoride, which is a gas.
Aluminium chloride, AlCl3, is a volatile solid which sublimes at 458K. The vapour formed on sublimation consists of an equilibrium mixture of monomers (AlCl3) and dimers (Al2Cl6). It is used to prepare the powerful and versatile reducing agent Lithium terahydridoaluminate, LiAlH4.
Both Boron chloride and Aluminium chloride act as Lewis acids to a wide range of electron pair donors, and this has led to their widespread use as catalysts. Aluminium chloride is used in the important Friedel-Crafts reaction.