C3.1 The Periodic Table

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  • Created by: StephBea
  • Created on: 31-03-15 20:44

1.1 The Early Periodic Table

  • The periodic table slowly evolved as chemists tried to classify the element, it arranges them in a pattern in which similar elements are grouped together.
  • During the 19th century many elements had been found, however scientists did not yet know the structure of atoms so decided to try and classify elements by their properties and weight.
  • Newland's table (1863) put the elements in order of atomic weight but failed to leave gaps for undiscovered elements, dispite the fact a new one was being discovered almost every year. He proposed the law of octaves, suggesting that every 8th element had similar properties. It was not accepted because although he put all 62 then discovered elements in his table, his theory only worked up until calcium.
  • Mendaleev's periodic table (1869) left gaps for the unknow elements and therefore provided a basis for our modern day periodic table. Mendaleev even predicted the properties of the missing elements, so when these elements were found and these predicted properties were confirmed Mandeleev's ideas became widly accepted by other scientists and it is upon his table that we base our modern day periodic table.
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1.2 The Modern Periodic Table

  • The atomic (proton) number determines the element's position on the table while the number of outer shells it has determines its chemical properties.
  • An element's group number (along the top of the table) is the same as the number of electrons each element in that column contains in its outer shell. As we go down a group the reactivity can be explained in terms of the distance between the outermost electron and the nucleus and the number of shells the atom contains.
  • An element's period number (down the side of the table) equals the number of outer shells (sometimes known as energy levels) that all the elements in that row contain.
  • An element's relative atomic mass can be seen in the top left hand corner of its box, this is the number of protons and neutrons within the atom. The number in the left bottom corner of the box is the atomic (proton) number. This is the number of protons in the element (equalling the number of electrons) as elements have no charge. 
  • The reactivity of metals as you go down a group increases, this is because metals lose electron when they react so the increasd number of shells means that the electrons in the outer shell are less strongly atracted to the nucleus because of distance and sheilding.
  • In contrast as you go down a group of non-metals the reactivity decreases. This is because when non-metals react they gain electrons, therefore the less shells there are, the stronger the attraction of electrons to the nucleus is.
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1.3 Group 1 - the Alkali Metals

  • All group 1 metals, when reacted with water, produce hydrogen gas and an alkaline solution containing the metal hydroxide. For example;
  • sodium + water --> sodium hydroxde + hydrogen / 2Na(s) + 2H2O(l) --> 2NaOH(aq) + H2(g)
  • Properties of the alkalis metals include; reacting readily with water and air, soft solids at room temperature, low melting points and low boiling points both of which decrease as you go down the group. All the elements in the group have low densities and the 3 elements with the lowest atomic masses; lithium, sodium and potassium, float on water. As with all groups of metals in the periodic table, the reactivity of the alkali metals increases as you go down.
  • Because the elements are group 1, all the elements all have just one electron in their outer shell, this electron is lost when they react to become stable (have a full outer shell) and this is why they all form positive ionic compounds of 1+ as they are left with one more proton than electron.
  • The alkali metals react with the halogens (group 7), to form salts that are white or colourless crystals, for example;
  • sodium + chlorine --> sodium chloride / 2Na(s) + Cl2(g) --> 2NaCl(s)
  • Compounds of alkali metals disolve in water, forming solutions that are normally colourless.
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1.4 The Transition Elements

  • The transition elements are found in the periodic table between groups 2 and 3, they are sometimes called transition metals because they are all metals. 
  • Compared with the akali metals, the transition elements have much higher melting and boiling points (except mercury) and densities. They are also much stronger and harder but are also less reactive. They do not react vigourously with either air or water.
  • The transition elements can form ions with different charges and compounds which are often coloured. We can find out the charge of a transition metal ion by looking at the Roman numberal number after its name. For example, iron(ii) chloride contains Fe2+ ions and so its formular is FeCl2, if it is iron(iii) then it contains Fe3+ ions so the formular would be FeCl3.
  • Transition elements and their brightly coloured compounds are important in industry for their use as catalysts.
  • Other properties of transition elements are that they are maleable (can be hammered or pressed into shape without breaking) and ductile (can be drawn into wires), they react very slowly or not al al with oxgyen or water at normal tempeatures and most are strong and dense making them great building materials, especially as alloys.
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1.5 Group 7 - the Halogens

  • The halogens all form ions with with a single negatively (1-) charged halide ion, in their ionic compounds with metals. They can also form covalent compounds by sharing electrons with other non-metals. 
  • A more reactive halogen can displace a less reactive halogen from a solution of one of its salts.
  • The reactivity of halogens, because they are non-metals, decreases as you go down the group because the have an almost complete outer shell so are gaining an electron to become stable and therefore the more shells they have then the more shielding taking place meaning the force of attraction to the neucleus is decreased.
  • All halogens exist as small molecules made up of a pair of atoms. Their properties include; low melting and boiling points that increase going down the group, at room temperature flourine is a yellow gas, chlorine is a green gas, bromine is a red-brown liquid and iodine is a grey solid. Iodine also easily evalpourates to a violet gas.
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