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Demonstrate understanding of aspects of selected elements

Demonstrate understanding of aspects of selected elements. Chemistry A.S. 1.4 2013. Periodic Table. We will concentrate on the following elements K, Na, Li, Mg, Ca, Al, Cu, Fe, Zn, Pb, Ag, Au, C, N, O, S, Cl, Br, I Elements can be classified as Metals or non-metals.

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Demonstrate understanding of aspects of selected elements

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  1. Demonstrate understanding of aspects of selected elements Chemistry A.S. 1.4 2013

  2. Periodic Table • We will concentrate on the following elements K, Na, Li, Mg, Ca, Al, Cu, Fe, Zn, Pb, Ag, Au, C, N, O, S, Cl, Br, I • Elements can be classified as Metals or non-metals. • Complete the exercise “structure of metals”

  3. Metals • Metals are found towards the left and bottom of the PT. • 1,2 or 3 valence electrons which are lost to form cations. • Metals form solids with metallic bonds. • The physical and chemical properties of metals are explained by the metallic solid structure.

  4. Physical properties of metals • Electrical conductivity • Thermal (heat) conductivity • Density (gml-1) • Lustre • Malleability • Colour • State • Ductility • MP and BP high high high >3 – except Na and Li high – when freshly cut high– more or less silver – except Cu and Au solid – except Hg High Vary but BP usually >1000oC The metals you need to know about and be able to relate their chemical and physical properties to their uses are… Na K Li Ca Mg Al Zn Fe Pb Cu Ag .

  5. Physical Properties Hardness – How easily a material can be cut or scratched with a knife. Many pure metals, like Iron, are too soft for engineering so mixtures are formed which increase their hardness – these are called alloys. “Alloys”

  6. Alloys • 2 or more molten metals are mixed and cooled to form a new mixture which has more useful properties than the parent metals. • When atoms of a different size are added the regular lattice structure is distorted so it is harder for layers of atoms to move, therefore the alloy is harder.

  7. Common Alloys

  8. Steel • Steel is an alloy made of metal and non-metal. (Fe and C) • It has high tensile strength. (can be bent and returned to its original shape without breaking) • Better corrosion resistance than Fe. • Other metals can be added to make specialist steels. Eg. Cr and Ni = stainless

  9. Uses of Metals “Wordsearch” / pg82 ESA

  10. Chemical Reactions of Metals • When metal atoms undergo chemical reactions new substances are formed. • All metals have low numbers of electrons in their valence shells so they have similar chemical properties. • Metal + Acid Metal Salt + Hydrogen • Hydrochloric acid forms CHLORIDE salts • Sulfuric acid forms SULFATE salts • Metal + Oxygen  Metal Oxide • Metal + Water  Hydrogen + Metal Hydroxide or Oxide • The ease of removal of electrons determines how reactive the metal is.

  11. Reactivity Series • Sodium • Lithium • Calcium • Magnesium • Aluminium • Zinc • Iron • Lead • Hydrogen • Copper • Silver • Gold Increases Chemical reactivity generally involves atoms gaining or losing electrons. Why?

  12. IONIC COMPOUNDS Metals lose valence electrons to form cations (+) Non metals gain valence electrons to form anions (-) The ions that have been formed are now attracted to each other. So Mg2+ will be attracted to Cl-. This forms an ionic compound. Naming rules: - The positive ion is first, and the negative second. - The negative ion ends in –ide, e.g. The sulfur atom becomes the sulfide ion. Exceptions: NO3- (nitrate) SO42- (sulfate) CO32- (carbonate) HCO3- (hydrogen carbonate)

  13. Ionic Bonding Strong electrical attractions (+/-) between oppositely charged ions in a 3D structure. • Na + Cl 2,8,1 2,8,7

  14. IONIC FORMULAE So Mg2+ will be attracted to Cl-. Because Mg is 2+ and Cl is only 1-, Mg will attract 2 Cl’s. The compound formed will be MgCl2. The subscript shows that the are 2 Cl’s for each Mg. If the starting ions were Cu2+ and S2-, the 2 ions have the same charge. So each Cu will only attract 1 S. The compound formed will be CuS. There is never any charges on the final product - they balance out

  15. DIFFICULT ONES NO3 Mg NO3 What is the formula for Magnesium Nitrate? Find the 2 ions on your table of ions… Mg2+ and NO3- This means that there are 3 oxygens attached to the Nitrogen – don’t let them get lost! So for each Mg we will need 2 NO3’s: The shorthand way of writing this is: Mg(NO3)2 The brackets are needed to show that we want 2 of the whole thing – you always need them if you have 2 of an ion with more than 1 bit. (polyatomic ions) NH4+ and S2- Fe3+ and OH- Al3+ and SO42- = (NH4)2S = Fe(OH)3 = Al2(SO4)3

  16. Metal reactions General Equation Word Equation • Metal + Oxygen → Metal oxide • Aluminium + Oxygen → Aluminium oxide • 4Al + 3O2 → 2Al2O3 • Metal + Water → Metal oxide + Hydrogen • Aluminium + Water → Aluminium oxide + Hydrogen • 2Al + 3H2O(g) → Al2O3 + 3H2 • Metal + Acid → Metal salt + Hydrogen • Magnesium + Hydrochloric acid → Magnesium Chloride + Hydrogen • Mg + 2HCl → MgCl2 + H2 Formula Equation

  17. Complete the following reactions: • Lithium + water • Lithium + hydrochloric acid • Silver + oxygen • Magnesium + sulphuric acid • Copper + oxygen • Aluminium + oxygen • Zinc + water (g) • Sodium + sulphuric acid • Lithium + oxygen • Aluminium + hydrochloric acid Lithium hydroxide + hydrogen Lithium chloride + hydrogen Silver oxide Magnesium sulphate + hydrogen Copper oxide Aluminium oxide Zinc oxide + hydrogen Sodium sulphate + hydrogen Lithium oxide Aluminium chloride + hydrogen

  18. METAL REACTIONS Metal + Oxygen Metal Oxide 1. METAL + OXYGEN Metals react in air to give metal oxides. Heating increases the rate of this reaction. (it may burn or change colour) Metal oxides are Basic, but only the first 2 groups of the periodic table are alkalis (bases that dissolve in water). An example: Magnesium is reacted in the air (with O2) to produce a white powder, which turns litmus paper blue. Write the word and balanced symbol equation for the reaction. Magnesium + Oxygen Magnesium oxide 2 Mg + O2 2 MgO General eqn: Ionic bond

  19. Metal + Water H2 + a Hydroxide 2. METAL + WATER Some metals react in water to give Hydrogen gas (H2) and a Metal Hydroxide or oxide. Reaction speed depends on the reactivity of the metal. Reactive metals react with cold water, others need steam . An example: Sodium reacts violently when placed in cold water and the gas produced sometimes explodes, but the reaction of Magnesium is only visible with steam and produces an oxide. Group 1 metal reactions Sodium + Water H2 + Sodium Hydroxide 2 2 Na + H2O H2 + NaOH 2 General eqn:

  20. Metal + Acid H2 + a Salt 3. METAL + ACID Many metals react in acid to give Hydrogen gas (H2) and a metal salt. Reaction speed depends on the reactivity of the metal. These ones react with acids. An example: Magnesium fizzes when placed in a test tube with Hydrochloric acid. The gas produced explodes with a squeaky pop. Magnesium + Hydrochloric acid H2 + Magnesium Chloride 2 Mg + HCl H2 + MgCl2 General eqn: “Equations practice1”

  21. ALUMINIUM Aluminium is high on the reactivity series, but never seems to do anything. Why? Aluminium forms an oxide coating very quickly. Aluminum oxide is shiny and silver so it looks like the metal but it doesn’t react. That is why aluminium is used for many things even though it is reactive. This is worth remembering. Examiners love to ask about it.

  22. Non-Metals • Top and right of the PT • 4,5,6 and 7 valence electrons • Electrons are usually gained to form anions • Non-metals form ionic compounds with metals and covalent bonds with other non-metals.

  23. Physical properties of non-metals • Non-metals are poor conductors – except for graphite • Colour • State The non-metals you need to know about and be able to relate their chemical and physical properties to their uses are… C N2 O2 O3 S Cl2 Br2 I2

  24. Physical properties If 2 or more forms of the same element exist in the same state, but with different arrangements of atoms they are called allotropes. “Diamond Vs Graphite”

  25. Allotropes of Carbon 1) Diamond – very hard, doesn’t conduct electricity, very high melting point , used for jewellery, cutting tools 2) Graphite – soft, shiny, does conduct electricity, very high melting point, used for lubricant, pencils, electrodes 3) Fullerenes – high tensile, conducts electricity, high ductility, dissolves in oil used for nanotechnology (molecular sponge)

  26. Graphene • A single layer of graphite is called graphene. • It has the ability to conduct electricity. • It can be rolled to form nanotubes. • It’s thinness and conducting ability make it a “super material” • Used for medicine delivery, carbon sequestration, wires in smaller electrical circuits.

  27. Allotropes of Oxygen • An electric spark can convert small amounts of 3O2 2O3 • O3 in the lower atmosphere is a respiratory pollutant. • In the upper atmosphere it blocks uv radiation • It can be used to purify air and water by killing microbes Ozone hole

  28. Sulfur • Yellow brittle crystal found near volcanoes Reactions • Burns with a blue flame • S + O2 SO2 Uses • “Sulfuric acid production”- contact process • Manufacture of fertiliser, medicine, fabric, explosives

  29. Contact Process • S + O2 SO2 (sulfur is burnt) • 2 SO2 + O2 2SO3 (lowers activation energy) • SO3 + H2SO4 H2S2O7 (dissolved to make oleum) • H2S2O7 + H2O 2 H2SO4 (gives off heat) 4000C, vanadium pentoxide Read and take notes on pg 96,97 ESA study guide, properties of Sulfuric acid

  30. Sulfur Dioxide Environmental effects • Acid Rain • Respiratory problems(choking, sharp smell) • S(s) + O2(g) SO2(g) • White smokey fumes, burns with a blue flame • Dissolves a bit in water • SO2 + H2O H2SO3 Sulfurous acid Uses • Bleach (oxidant) • Preservative of fruit • winemaking

  31. Chlorine • Found as NaCl and combined with other metals. • Cl2 is poisonous yellow/green gas Use • Kills microbes in water • Cl2 + H2O HCl + HOCl • HOCl H+ + OCl- (hypochlorite ion oxidises/ kills bacteria) • Household bleach • 2NaOH + Cl2NaOCl + NaCl + H2O (hypochlorite ion sterilises) • Monomer component of PVC • Bleaching agent in the paper industry Makes water acidic

  32. Bromine • Found combined with other metals egNaBr • Br2is poisonous red/brown liquid Use • Kills microbes in water • Medications • Light sensitive agent in photography

  33. Iodine • Blue / Black solid at room temperature • Sublimes to form violet/ pink gas • Very rare but found as a soluble iodide salt in seawater and many seaweeds. Use • Medicines, dyes, catalysts, essential in diet

  34. Trends down group 17 • Atoms in the same group have the same number of valence electrons and therefore react similarly. • Atoms higher in the group have less “shells” so the valence electrons are closer to the positively charged protons in the nucleus. • The smaller atoms therefore have greater attraction for electrons and are more reactive.

  35. Nitrogen • Colourless, odourless atmospheric gas • Sodium nitrate (saltpetre) found in Earth’s crust Reactions • N2 + O2 2NO • 2NO + O2 2NO2 Ammonia, NH3 is a strong smelling gas which is water soluble. • NH3 + H2O NH4+ + OH- • Ammonia is a base and is therefore neutralised to make ammonium compounds. • NH3 + HClNH4Cl Makes water basic

  36. Ammonia Uses • Making HNO3 (making explosives) • Liquid refrigerant • Agricultural supplement • Industrial “ammonia production” • Lab production • Ca(OH)2 + 2NH4Cl 2NH3 + 2H2O + CaCl2

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