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Aims

Aims. To increase awareness of chemical formula and structure. To establish a concrete knowledge of ionic and covalent bonding. To make students aware of the impact of electronegativity on bonding. To explore shapes of molecules and intermolecular forces.

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Aims

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  1. Aims • To increase awareness of chemical formula and structure. • To establish a concrete knowledge of ionic and covalent bonding. • To make students aware of the impact of electronegativity on bonding. • To explore shapes of molecules and intermolecular forces. • To explore the social and applied aspects of this topic.

  2. Learning Outcomes • Define covalent and ionic bonding. • Illustrate examples of covalent and ionic bonding. • Demonstrate that bonding involves valence electrons. • Identify from a formula the probable type of bond. • Discuss the differences between various intermolecular and intramolecular forces. • Draw a dot and cross diagram and structural formula for simple molecules and ions and deduce the molecular formula.

  3. Polarity in Molecules • Polar molecules have a centre of positive charge separated from the centre of the negative charge. Even though the individual bonds may be non- polar the overall molecule is polar for this reason. E.G. NH3 and H2O • Non-polar molecules have the centre of the positive and negative charges in one location. There is an equal pull on the e- even though the individual bonds may be polar. E.G. CCl4 and CO2

  4. Inter and Intra Molecular Forces • The force of attraction between ions is stronger than between molecules. • Inter: are forces between molecules. • Intra: are forces within a molecule. • There are 3 kinds of forces that can attract molecules together

  5. Van der Waal’s: • These are the weakest forces caused by the movement of e- within a molecule. The electrons move randomly within the bond so at 1 point in time they are nearer to 1 atom than the other. This induces a temporary dipole force. • Temporary dipoles will result in increased boiling points. • The greater number of e- in a molecule the greater the number of temporary dipoles.

  6. Dipole-dipole: • The positively charged end of a polar molecule is attracted to the negative end of another molecule. • The dipoles in this case are permanent. • As a result they are stronger than Van der Waal’s forces.

  7. Hydrogen Bonding • When H is bonded to F, O or N these elements are sufficiently electronegative to make the bond polar. • H has only 1 e- in its atom, a strong partial positive charge will result. • This means it is very strongly attracted to the negative atom and as a result H2O is a liquid at room temperature with a fairly high boiling point.

  8. Experiments on Bonding • Qualitative tests for the anions CO32-, HCO3-, SO42-, SO32-, Cl-, NO3- and PO43- in aqueous solution. • Theory: Reactions of anions with certain reagents to produce characteristic coloured precipitates or other easily identifiable results are employed to identify or confirm the presence of these anions in aqueous solution and to distinguish the anions from one another.

  9. (a)To test for the carbonate and hydrogencarbonate anions. • NB: Wear your safety glasses. • 1. Add 2 cm3 of sodium carbonate solution to a clean test tube labelled A, and 2 cm3 of sodium hydrogencarbonate solution to a clean test tube labelled B. • 2. Using a dropping pipette add 2 cm3 of dilute hydrochloric acid to each test tube. Record your observations.

  10. Add 2 cm3 of sodium carbonate solution to a clean test tube labelled C, and 2 cm3 of sodium hydrogencarbonate solution to a clean test tube labelled D. Using a dropping pipette add 2 cm3 of magnesium sulphate solution to each. Record your observations. • Carefully heat the contents of test tubes labelled C and D at the end of the last step.

  11. (b)To test for the sulphate and sulphite anions • Add 2 cm3 of sodium sulfate solution to a clean test tube labelled A and 2 cm3 of sodium sulfite solution to a clean test tube labelled B. • Using a dropping pipette add 1 cm3 of barium chloride solution to each. Record your observations.

  12. Add 2 cm3 of dilute hydrochloric acid to the contents of the test tubes labelled A and B at the end of the last step. Record your observations.

  13. (c)To test for the chloride anion • Add 2 cm3 of sodium chloride solution to a clean test tube. Using a dropping pipette add a few drops of silver nitrate solution. Record your observations • Add 2 cm3 of dilute ammonia solution to the contents of the test tubes at the end.

  14. (d)To test for the nitrate anion • Add 2 cm3 of potassium nitrate solution to a clean test tube. Using a dropping pipette, add 3 cm3 of cold saturated iron(II) sulfate solution. • Using a dropping pipette, carefully add 2 cm3 of concentrated sulfuric acid slowly down the wall of the test tube. Do not mix the contents of the test tube. Record your observations.

  15. (e)To test for the phosphate anion • Add 2 cm3 of disodium hydrogenphosphate(V) solution to a clean test tube. • Using a dropping pipette add approximately 6 cm3 of the clear ammonium molybdate reagent to the test tube. • Warm gently by placing in a beaker of water at a temperature not exceeding 40 C. Record any observations. • Add an equal volume of ammonia solution to the contents of the test tube at the end.

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