Periodic table 2013

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  • 1. Periodic Table of Elements

2. Analyse the Periodic Table of Elements Analyse Group 18 elements Analyse Group 1 elements Analyse Group 17 elements Analyse elements in a period Understand transition elements 3. Antoine Lavoisier (1743 1794) Classify substances into metals and non-metals Unsuccessful because light, heat and some other compounds where not elements. 4. Johann Dobereiner (1780 - 1849) Introduced triads. Elements were classified into groups of three elements with same chemical properties The atomic mass of middle elements was approximately the average atomic mass of the other two elements 5. Lothar Meyer (1830 - 1895) Plotted a graph of the atomic volume against atomic mass. Elements with similar chemical properties occupied same positions. Successful in showing the properties of elements formed a periodic pattern against their atomic masses. 6. John Newlands (1837 - 1898) Arranged elements in order of increasing atomic mass. Elements with similar properties recurred at every eight element. This was known as the Law of Octaves Failed because only obeyed by first 17 elements only 7. Dimetri Mendeleev (1839 1907) Arranged elements in order of increasing atomic mass Elements with similar chemical properties are grouped together He left empty spaces in the table for undiscovered elements 8. Henry J. G. Moseley (1887-1915) Concluded that proton number should be the bases for the periodic change of chemical properties Arranged the elements in order of increasing proton number in the Periodic Table. 9. Elements are arranged according their increasing order of proton number. Vertical columns = groups(according to their number of valence electrons) Horizontal rows = periods (number of electron shells filled by electrons) 10. Known as noble gases/inert gases(chemically unreactive elements) Non-metals that exist as monoatomic colourless gases. Members : Helium(He) Neon(Ne) Argon(Ar) Krypton(Kr) Xenon(Xe) Radon(Rn). 11. Very small atomic sizes. Low melting and boiling points Weak van der Waals forces of attraction between atoms. Low densities Very small masses but huge volumes. Melting and boiling points of elements increase down the Group 18. 12. All Group 18 elements are chemically inert/unreactive. The outermost electron shell of each member is fully occupied by electrons. This is a stable electron arrangement which in Helium, it is said achieve duplet electron arrangement. Other than Helium, it is said achieve octet electron arrangement. 13. Helium To fill airships and weather balloons. used as artificial atmosphere in oxygen tank for divers. 14. Neon Advertising lights Used in aeroplane runway lights 15. Argon To fill light bulbs. Krypton Used in lasers to repair the retina of the eye. To fill photographic flash lamps. 16. Xenon Making electron tubes and stroboscopic lamps Used in bubble chambers in atomic energy reactors. Radon Used to treat cancer 17. Consists of lithium(Li), sodium(Na), potassium(K), rubidium(Rb), caesium(Cs) and francium(Fr). Li Na K Rb Cs Fr They are known as alkali metals because they react with water to produce alkaline solution. 18. Soft Low melting points Low densities Shiny and silvery surface Good conductor of heat Good conductor of electricity 19. Hardness, melting point and boiling point of the elements decreases going down the group. When go down Group 1, size of atom becomes larger. The positive nucleus gets further away from the negative sea of electrons. The force of attraction between the metal ions and the sea of electrons gets weaker down the group. Less energy is needed to overcome this weakening force of attraction. 20. 1. All react with water to produce alkaline metal hydroxide solution and hydrogen gas. 2X(s) + 2H2O(l) 2XOH(aq) + H2(g) 21. 2. All burn in oxygen gas to produce white solid metal oxides. 4X(s) + O2(g) 2X2O(s) The oxide dissolve in water to form alkaline metal hydroxide solution. X2O(s) + H2O(l) 2XOH(aq) 3. All burn in chlorine gas to produce white solid metal chlorides. 2X(s) + Cl2(g) 2XCl(s) 22. Why the reactivity of elements increases down the Group 1? Atomic size of Group 1 elements increases from lithium to francium//Number of shells occupied by electrons increases. Distance between the valence electron in the outermost shell and positive nucleus increases down the Group 1. Attraction between nucleus and valence electron decreases. It is easier for the atom to lose the valence electron to achieve stable electron arrangement. Why all elements in Group 1 have same chemical properties? Chemical reaction is all about the activity of electrons All the elements have one valence electron. Each of them reacts by donating one valence electron to form an ion with a charge of +1 to achieve stable electron arrangement. 23. Members are fluorine(F2) , chlorine(Cl2), bromine(Br2), iodine(I2), and astatine(At2) F Cl Br I At The elements are also known as halogens which exist as diatomic molecules. 24. They have low melting and boiling points because molecules are attracted to each other by weak van der Waals forces of attraction. The melting and boiling points of the elements increases down Group 17. This change the states of elements from gas to solid and the colour of elements from lighter colour to darker colour. 25. Elements State Colour Fluorine Gas Pale yellow Chlorine Gas Greenish-yellow Bromine Liquid Reddish-brown Iodine Solid Purplish-black 26. Why the melting and boiling points of elements increases down Group 17? Molecular size/relative molecular mass of the elements increases down Group 17. Forces of attraction between molecules/Intermolecular forces of attraction increases. More heat is needed to overcome the stronger forces of attraction between the molecules. 27. All members have similar chemical properties but differ in the reactivity. 1. React with water to form two acids X2(g) + H2O(l) HX(aq) + HOX(aq) Example: Cl2(g) + H2O(l) HCl(aq) + HOCl(aq) hydrochloric hypochlorous acid acid Hypochlorous acid is a bleaching agent (bleach both blue and red litmus paper) 28. 2. Halogens in gaseous state react with hot iron to form brown solid. 2Fe(s) + 3X2(g) 2FeX3(s) Example: 2Fe(s) + 3Cl2(g) 2FeCl3(s) solid iron(III) chloride(brown) 3. Halogens react with sodium hydroxide solution to produce sodium halide, sodium halate(I) and water X2 + 2NaOH(aq) NaX(aq) + NaOX(aq) + H2O(l) Example: Cl2 + 2NaOH(aq) NaCl(aq) + NaOCl(aq) + H2O(l) Sodium chlorate(I) 29. Why all halogens possess similar chemical properties? Chemical reaction = lose or accept electrons All halogens always gain one electron to achieve stable octet electron arrangement. Therefore, they have similar chemical properties. Why chemical reactivity of halogens decreases down Group 17? Atomic size/number of electron occupied shells of halogens increases down Group 17. The outermost shell becomes further from the nucleus of the atom. Strength to attract one electron into the outermost shell by the nucleus becomes weaker. Reactivity decreases. 30. Elements across a period exhibit a periodic change in properties. Proton number increases by one unit from one element to the next element 31. All the atoms of the elements have three shells occupied with electrons The number of valence electrons in each atom increase from 1 to 8 32. All the elements exist as solid except chlorine and argon which are gases The atomic radius of elements decreases. This is due to the increasing nuclei attraction on the valence electrons. The electronegativity of elements increases. This is also due to the increasing nuclei attraction on the valence electrons and the decreases in atomic size. 33. Uses of metalloid Make diodes and transistors A diode A transistor Both are commonly used in the making of microchips Microchips are widely used in the manufacture of computers, mobile phones, televisions, video recorders, calculators, radio and etc. Metalloid semi-metal, reacts with acid only, weak conductor, brittle and not malleable and ductile. 34. Oxides of elements change from basic to amphoteric and then to acidic across the period towards the right. Basic oxides react with acids to form salt and water Acidic oxides react with alkalis to form salt and water Amphoteric oxides react with both acids and alkalis to form salt and water. 35. Elements from Group 3 to Group 12 in the Periodic Table. Common characteristics Solid metal with shiny surface. Good conductor of heat and electricity. High melting and boiling points. Hard, malleable and ductile. 36. Special characteristics; Show different oxidation numbers in their compounds Form coloured ions or compounds Use as catalysts Form complex ions 37. Show different oxidation numbers in their compound Compound Formula Oxidation number Chromium(III) chloride CrCl3 +3 Potassium dichromate(VII) K2Cr2O7 +6 Manganese(II) sulphate MnSO4 +2 Manganese(VI) oxide MnO2 +4 Potassium manganate(VII) KMnO4 +7 Iron(II) sulphate FeSO4 +2 Iron(III) chloride FeCl3 +3 Copper(I) oxide Cu2O +1 Copper(II) sulphate CuSO4 +2 38. Form coloured ions or compounds Element Ion Colour Chromium Cr3+ Green CrO4 2- Yellow Cr2O7 2- Orange Manganese Mn2+ Pale pink MnO4 - Purple Iron Fe2+ Pale green Fe3+ Yellowish brown Cobalt Co2+ Pink Nickel Ni2+ Green Copper Cu2+ Blue Green 39. Form coloured ions or compounds Gemstone Transition metal Colour Emerald Ni and Fe Green Amethyst Fe and Mn Purple Sapphire Co and Ti Blue Ruby Cr Red Topaz Fe Yellow 40. As catalyst Process Catalyst To manufacture Haber Process Iron fillings, Fe Ammonia Contact Process Vanadium(V) oxide, V2O5 Sulphuric acid Ostwald Process Platinum, Pt Nitric acid Hydrogenation Nickel, Ni Margarine HAI CSV ONiP 41. To form complex ions Element Complex ions Formula Iron Hexacyanoferrate(II) ion [Fe(CN)6]4- Hexacyanoferrate(III) ion [Fe(CN)6]3- Chromium Hexaamina chromium(III) ion [Cr(NH3)6]3+ Copper Tetraami